7 

i  II ii  ii 


,  PP 


PLATE  OF  SPECTRA 

The  numbers  at  top  and  bottom  of  plate  give  wave  lengths 
for  each  spectrum,  in  hundredths  of  a  micron  (//). 


TEXTBOOK  OF  CHEMISTEY 


BY 
WILLIAM   A.   NOYES 

DIRECTOR   OF   THE    CHEMICAL   LABORATORY 
OF   THE    UNIVERSITY   OF    ILLINOIS 


NEW  YORK 

HENRY   HOLT  AND   COMPANY 
1919 


iv  PREFACE 

time  displaced  by  something  very  different.  It  is  believed 
that  the  development  of  our  knowledge  during  the  last  few 
years  fully  justifies  this  course.  The  theory  of  ionization  has 
also  been  freely  used,  as  the  only  means  we  have  by  which 
a  large  class  of  phenomena  can  be  clearly  presented  and 
understood. 

It  seems  desirable  to  give  some  material  which  it  is  not 
possible  to  emphasize  or  teach  thoroughly  in  a  brief  course 
and  some  things  which  are  rather  for  reference  than  to  be 
learned.  To  aid  teachers  and  students  in  distinguishing  such 
paragraphs,  they  are  indicated  by  an  asterisk.  This  device, 
which  was,  I  think,  first  used  by  Professor  Ostwald,  is  better 
adapted  than  the  use  of  fine  print  to  the  need  of  teachers  who 
may  wish  to  make  a  different  selection.  Students  are  earnestly 
advised  to  read  these  paragraphs,  even  when  they  are  not 
expected  to  acquire  a  full  knowledge  of  them. 

I  wish  to  express  my  very  sincere  thanks  to  the  following 
members  of  the  Chemical  Staff  of  the  University  of  Illinois, 
some  of  whom  have  met  with  me  almost  weekly  for  three 
years  to  read  and  criticize  the  successive  chapters  of  the  book. 
The  criticisms  have  been  very  valuable  and  helpful.  C.  W. 
Balke,  S.  J.  Bates,  Edward  Bartow,  G.  D.  Beal,  L.  L.  Burgess, 
E.  S.  Curtiss,  C.  G.  Derick,  Lambert  Thorp,  B.  S.  Hopkins, 
Helen  Isham  (Mattill),  Grinnell  Jones,  C.  G.  MacArthur,  Ellen 
S.  McCarthy  (Foley),  D.  F.  McFarland,  D.  A.  Maclnnes,  C.  F. 
Nelson,  S.  W.  Parr,  G.  McP.  Smith,  E.  K.  Strachan,  E.  W. 
Washburn,  H.  C.  P.  Weber. 

I  wish  also  to  express  my  gratitude  to  Professor  Edward  W. 
Morley,  Professor  Julius  Stieglitz  of  the  University  of  Chi- 
cago, Professor  J.  Bishop  Tingle  of  MacMaster  University, 
and  Mr.  C.  M.  Wirick  of  the  Crane  Technical  High  School  of 
Chicago,  who  have  read  the  proofs  and  made  many  useful  sug- 
gestions. I  am  also  indebted  to  Professor  K.  B.  Moore  of  the 
Bureau  of  Mines  for  some  valuable  criticisms,  of  the  paragraph 
on  Radioactivity. 


CONTENTS 

CHAPTER  I 
INTRODUCTION 

The  Nature  of  Scientific  Knowledge,  1.  —  Subdivisions  of  Science,  3.— 
Physical  Sciences,  4.  —  Matter  and  Energy,  5. — Conservation  of  Matter  and 
Energy,  6.  —  Pure  Substances  and  Mixtures,  7. —  Preparation  of  Pure  Sub- 
stances, 8.  —  Elements  and  Compounds,  9.  —  How  Pure  Substances  are  dis- 
tinguished from  Mixtures.  Law  of  Constant  Proportion,  12.  —  Inductive 
Reasoning,  13.  — Law  of  Combining  Weights,  13. —The  Atomic  Theory,  14. 
—  Selection  of  Atomic  Weights,  16.  —  Formulas,  16.  —  Composition  of  Pure 
Substances,  17.  —  Study  of  Chemistry,  18. 


CHAPTER  II 

OXYGEN 
SYMBOL,  O.    ATOMIC  WEIGHT,  16 

Occurrence,  19.— Preparation,  19.  —  Collection  and  Storage  of  Gases,  22.— 
Properties  of  Oxygen,  22.  —  Oxygen  and  Acid  Properties,  23.  —  Combustion. 
Effect  of  Concentration  on  a  Chemical  Reaction,  24.  —  Kindling  Tempera- 
ture, 24.  — Heat  of  Combustion.  Calorimeter,  25.  — The  Nature  of  Chemical 
Energy,  27.  — Catalysis,  28.  — Chemical  Affinity,  29.  — Nomenclature,  29. 


CHAPTER  III 

LAWS  OF  GASES 
UNITS  OF  LENGTH,  WEIGHT,  VOLUME,  TEMPERATURE,  TIME  AND  ENERGY 

Unit  of  Length.  Meter,  31.  — Unit  of  Weight.  Gram,  31.  — Unit  of  Vol- 
ume. Liter,  31.— Units  of  Time,  32.  — Unit  of  Temperature,  32.  — Units  of 
Energy.  Kilogram-meter.  Erg,  32.  —  Centimeter-gram-second  System.  Ab- 
solute Units,  33.  — Units  of  Mechanical  Energy,  33.— Unit  of  Power,  33.— 
Units  of  Heat,  33.  —  Electrical  Units,  33.  — Chemical  Energy,  34.  — Effect  of 
Pressure  on  a  Gas.  Law  of  Boyle,  34. — Corrections  for  Readings  of  the 
Barometer,  36.  — Effect  of  Temperature  on  a  Gas.  Law  of  Charles,  38.— 
Absolute  Temperatures,  39.  —  Significance  of  the  Absolute  Zero,  40. — Deter- 
mination of  the  Weight  of  a  Liter  of  a  Gas,  40.  —  Graphical  Representation 
of  the  Gas  Laws,  42.— Exercises,  43. 

v 


yi  CONTENTS 

CHAPTER  IV 

HYDROGEN 
SYMBOL,  H.    ATOMIC  WEIGHT,  1.0078 

Occurrence,  45.— Radicals,  46.— Salts,  47.  — Preparation  of  Hydrogen. 
1,  Electrolysis  of  Dilute  Sulf uric  Acid,  47.— Electrolytes.  Ions.  Theory  of 
Electrolysis,  48.—  2,  Preparation  of  Hydrogen  from  Iron  and  Steam,  48. 
Reversible  Reactions,  50.—  3,  Decomposition  of  Water  by  Metals  at  Ordi- 
nary Temperatures,  50. — Contrast  between  the  Action  of  Iron  and  of  Sodium 
on  Water,  51. — 4,  Hydrogen  from  "Hydrone,"  52.  —  5,  Preparation  of 
Hydrogen  by  the  Action  of  Metals  on  Acids,  52.  —  Apparatus  for  the  Prepa- 
ration of  Hydrogen,  53.  —  Purification  of  Hydrogen,  54.  —  Properties  of 
Hydrogen,  55.  — Diffusion  of  Gases,  56.— Kinetic  Theory  of  Gases,  58.— 
Chemical  Properties  of  Hydrogen,  59.  — Dissociation,  59.— The  Oxyhydrogen 
Blowpipe,  61.— Explosions.  Catalysis,  62.  — Oxidation.  Reduction,  63.— 
Valence,  63.  — Heat  of  Combustion  of  Hydrogen,  65. 

CHAPTER   V 
WATER,  HYDROGEN  PEROXIDE 

Analysis,  Synthesis,  66.  —  Qualitative  Analysis  and  Synthesis  of  Water,  66. 

—  Quantitative  Synthesis  of  Water  by  Volume,  66.  —  Composition  of  Water 
by  Weight,  68.— The  Unit  for  Atomic  Weights,  68.  — Determination  of  the 
Composition  of  Water  by  the  Use  of  Copper  Oxide,  69.  —  Determination  of 
the  Composition  of  Water  by  Weighing  Oxygen  and  Hydrogen,  71.  — Proper- 
ties of  Water,  72.  —Heat  of  Fusion  and  Vaporization,  74.  —  Vapor  Pressure  of 
Water,  74.  — Equilibrium,  76.  — Effect  of  Water  Vapor  on  the  Pressure  of  a 
Gas,  76.— Phases.    Degrees  of  Freedom,  77.  — Water  as  a  Solvent.    Solutes, 
79.— Chemical  Activity  in  Solutions.    Metathesis,  81.  —  Hydrates,  Deliques- 
cence, Efflorescence,  81.  — Natural  Waters,  82.  — Purification  of  Water,  83.— 
Hydrogen  Peroxide,  83.  —  Properties  and  Uses  of  Hydrogen  Peroxide,  86. — 
Tests  for  Hydrogen  Peroxide,  85.  —  Structure  of  Hydrogen  Peroxide,   86. 

—  Law  of  Multiple  Proportion,  87. 

CHAPTER   VI 
AVOGADRO'S  LAW.    SELECTION  OF  ATOMIC  WEIGHTS.    OZONE 

Gay  Lussac's  Law  of  Combining  Volumes,  89.  —  Avogadro's  Law,  91.  — 
Selection  of  an  Atomic  Weight,  92.  —  Molecules  of  the  Elements,  93.  —  Gram 
Molecular  Volume,  94.  —  Number  of  Molecules  in  one  Cubic  Centimeter  of  a 
Gas,  95.  —  Allotropic  Forms.  Ozone,  97.  —  Exercises,  99. 

CHAPTER   VII 

CHLORINE 
SYMBOL,  CL.    ATOMIC  WEIGHT,  35.46.    FORMULA,  C12 

Occurrences  of  Chlorine,  100.  —  Preparation  of  Chlorine.  1,  By  Electrol- 
ysis of  Sodium  Chloride,  100.  —  2,  Preparation  by  Oxidation  of  Hydrochloric 


CONTENTS  Vii 

Acid,  100.  —  3,  Preparation  of  Chlorine  by  the  Deacon  Process,  102.  —  4,  The 
Weldon  Process  for  Chlorine,  103.  —  Properties  of  Chlorine,  104.  —  Chlorine 
and  Water.  Bleaching,  106.  —  Chlorine  Hydrate.  Phases,  107. —The  Heat 
of  Combination  of  Chlorine  and  of  Oxygen  with  Other  Elements,  108.  —  Equi- 
librium in  Chemical  Reactions,  108.  —  Principle  of  van't  Hoff-Le  Chatelier, 
111.  —Effect  of  Water  on  Chlorides.  lonization,  112.  —Effect  of  Water  on 
Chlorides.  Hydrolysis,  115.  —  Exercises,  116. 


CHAPTER   VIII 

HYDROCHLORIC  ACID.    OXIDES  AND  OXYACIDS  OF  CHLORINE 

Hydrochloric  Acid,  118.  —  Properties  of  Hydrochloric  Acid,  119.  —1,  Reac- 
tion with  Metals,  120.  —2,  Reaction  with  Hydroxides  of  Metals,  121.  —3,  Re- 
action with  Oxides  of  Metals,  122.  —  4,  Reaction  with  Oxidizing  Agents,  122. 
—  Indicators,  122.  —  Oxides  and  Oxygen  Acids  of  Chlorine.  Nomenclature, 
123.  —  Hypochlorous  Acid.  Hypochlorites,  124.  —  Hypochlorous  Anhydride, 
or  Chlorine  Monoxide,  126.  —  Chlorous  Acid  and  Chlorites,  127. —  Chloric 
Acid  and  Chlorates,  127.  —  Chlorine  Peroxide,  127.  —  Perchlorates  and  Per- 
chloric Acid,  128.  —  Structure  of  the  Oxyacids  of  Chlorine,  130.  —  The 
Atomic  Weight  of  Chlorine,  130. 

CHAPTER  IX 
CLASSIFICATION  OF  THE  ELEMENTS.    THE  PERIODIC  SYSTEM,  132 

CHAPTER  X 
THE  HALOGEN  FAMILY 

General  Properties  of  the  Halogens,  139.  —  Compounds  of  the  Halogens 
with  Hydrogen  and  Oxygen,  138.  —  Bromine,  Br,  79.92.  Occurrence.  Prepa- 
ration, 140.  —  Properties,  141.  —  Hydrobromic  Acid,  142. —  Sodium  Hypo- 
bromite,  143.  Iodine,  I,  126.92.  Occurrence,  Preparation,  144.  —  Properties 
of  Iodine,  144.  —  Hydriodic  Acid,  145.  —  Direct  Combination  of  Hydrogen  and 
Iodine.  Reversible  Reactions.  Equilibrium,  146.  —  Speed  of  Chemical  Reac- 
tion, 148.  —  Concentration  and  Speed  of  Reaction,  149.  —  Calculation  of  the 
Relative  Speed  of  Two  Reactions  from  the  Composition  of  an  Equilibrium 
Mixture,  151.  — Effect  of  Removing  One  of  the  Reacting  Substances.  Dis- 
placement of  the  Equilibrium  Point,  152.  —  Heat  of  Formation  of  Hydriodic 
Acid,  152.  —  Flourine,  F,  19.0.  Occurrence,  153.  —  Preparation,  153.  —  Prop- 
erties, 154.  —  Etching  Glass.  Hydrofluoric  Acid,  154.  —  Metallic  Elements  of 
Group  VII,  156.  —  Exercises,  156. 

CHAPTER  XI 

SULFUR,   SELENIUM  AND  TELLURIUM 

Sulfur,  S,  32.0.  Occurrence,  160.  —  Allotropic  Forms  of  Sulfur,  162.— 
Liquid  Forms  of  Sulfur,  162.  —  Gaseous  Forms  of  Sulfur,  163.  —  Properties 


viii  CONTENTS 

and  Uses  of  Sulfur,  163.  —  Hydrogen  Sulfide,  164.  —  Solution  of  Hydrogen 
Sulfide.    Henry's  Law,  165.  —  Sulfides.    Groups  of  Analytical  Chemistry,  166. 

—  Hydrosulfuric  Acid.     Strength  of  Acids,  167.  —  Application  of  the  Idea  of 
Strength  of  Acids  to  Explain  the  Conduct  of  Sulfides,  168.  —  Hydrogen  Sul- 
fide as  a  Reducing  Agent,  171.  —  Sulfur  Dioxide,  172.  —  Sulfurous  Acid,  174. 

—  Sulfites,  175.  — Sulfur  Trioxide,  175.  —  Sulf uric  Acid,  177.— The  Electron 
Theory,  181.  —  Sulf  uric  Acid  as  a  Dehydrating  Agent,  182.  —  Sulfates.    Di- 
basic Acids,  183.  —  Normal,  Standard  and  Formular  Solutions,  183.  —  Acid- 
imetry    and    Alkalimetry,    185.  —  Pyrosulfates,    186.  —  Hyposulfites,    186.— 
Thiosulfates,  186.  —  Persulfuric  Acid,   187.  —  Permonosulfuric  Acid,   188.— 
Polythionic  Acids,    188.  —  Compounds  of  Sulfur  Containing  Halogens,  188. 

—  Sulfur  Monochloride,  188.  -  Chlorosulfonic  Acid,  189.  —  Sulfuryl  Chloride, 

189.  —  Selenium,  Se,  79.2,  189.  —  Hydrogen  Selenide,  190.  —  Selenium  Dioxide, 

190.  —  Selenic  Acid,  190.  —  Tellurium,  Te,  127.5,  190.  — Atomic  Weight  of 
Tellurium,  190.  —  General  Properties  of  the  Elements  of  the  Sixth  Group, 

191.  —  Crystals,  192.  —1,  The  Isometric  or  Regular  System,  193.  —  2,  The 
Tetragonal  System,  194.  —  3,  The  Rhombic  System,  194.—  4,  The  Hexagonal 
System,  194.  — 5,  The  Monoclinic  System,  195.  —  6,  The  Triclinic  System,  196. 

—  Exercises,  196. 


CHAPTER  XII 

NITROGEN 
SYMBOL,  H.    ATOMIC  WEIGHT,  14.01 

Occurrence  and  Natural  History  of  Nitrogen,  198.  —  Preparation  and  Prop- 
erties of  Nitrogen,  200.  —  Ammonia,  201.  —  Properties  of  Ammonia,  202.  — 
Aqua  Ammonia,  203.  —  Ice  Machines,  204.  —  Derivatives  of  Ammonia,  205.  — 
The  Electron  Theory,  206.  —  Solutions  in  Liquid  Ammonia,  207.  — The  Volu- 
metric Composition  of  Ammonia,  208.  —  Nitric  Acid,  210.  —  Hydrates  of  Nitric 
Acid,  211.  —  Chemical  Properties  of  Nitric  Acid,  212.— Aqua  Regia,  213.— 
Nitrosyl  Chloride,  214.  — Oxides  of  Nitrogen,  214.  — Nitrous  Oxide,  214.— 
Nitric  Oxide,  215.  —  Nitrous  Anhydride,  218.  —  Nitrous  Acid,  218.  —  Nitrogen 
Peroxide  and  Nitrogen  Tetroxide,  219.  —  Nitrogen  Pentoxide,  220. —Other 
Compounds  of  Nitrogen,  220.  —  Hyponitrous  Acid,  221.  — Hydro xylamine,  221. 
—  Hydrazine,  222.  —  Hydronitric  Acid,  or  Azoimide,  223.  —Iodine  Trinitride, 
223.  —  Nitrogen  Trichloride,  224.  —  Nitrogen  Iodide,  225.  — Nitro  Nitrogen 
Trichloride,  225.  —  Endothermic  Compounds,  225.  —  Exercises,  225. 


CHAPTER  XIII 

THE  ATMOSPHERE.  NOBLE  GASES 

Determination  of  Oxygen,  227.  —  Composition  of  Air,  227.  — Air  is  a  Mix- 
ture, 228. —Carbon  Dioxide  in  the  Air,  229.  — Ventilation,  230.  —  Moisture, 
231.  — Liquid  Air.  Critical  Temperature,  232.  — Argon,  A,  39.88,  235.— 
Atomic  Weight  of  Argon.  Specific  Heat  of  Gases,  236. —  Helium,  He,  3.99, 
237.  — Neon,  Krypton,  Xenon  and  Niton,  238.  —  Exercises,  238. 


CONTENTS  ix 

CHAPTER  XIV 
PHOSPHORUS 

Phosphorus  ,  P,  31.04.    Occurrence,  240.  —  Preparation  of  Phosphorus,  241. 

—  Allotropic  Forms  of  Phosphorus,  241.  —  Matches,  242.  —  Phosphine,  243.  — 
Phosphonium  Salts,  243.  —  Phosphorus  Trichloride,  and  Phosphorus  Penta- 
chloride,  244.  —  Hydrolysis  of  the  Chlorides  of  Phosphorus,  245.  —  Phosphorus 
Oxychloride,  245.  —  Oxides  of  Phosphorus,  246.  —  Acids  of  Phosphorus,  246.  — 
Basicity  of  the  Acids  of  Phosphorus,  247.  —  Hypophosphorous  Acid,  248. — 
Phosphorous  Acid,  248.  —  Orthophosphoric  Acid,  248.  —  lonization  of  Ortho- 
phosphoric  Acid,  250.  —  Decomposition  of  Primary  and  Secondary  Salts  of 
Orthophosphoric  Acid,  252.  —  Pyrophosphoric   Acid,  253.  —  Metaphosphoric 
Acid,   253.  —  Hypophosphoric   Acid,    254.  —  Sulfides    of    Phosphorus,  254.— 
Exercises,  255. 

CHAPTER  XV 
ARSENIC,  ANTIMONY  AND  BISMUTH 

Arsenic,  As,  74.96.  Occurrence,  256.  —  Preparation  and  Properties  of  Arsenic, 
257.  —  Arsine,  Marsh's  Test,  257.  —  Arsenic  "Trioxide,"  258.  —  Arsenious  Acid, 
259.  —  Arsenic  Pentoxide  and  Arsenic  Acid,  259.  —  Arsenic  Trichloride,  260. 

—  Sulfides  of  Arsenic,  260.  — Arsenic  Disulfide,  or  Realgar,  260.  — Arsenic 
Trisulfide,  or  Orpiment,  260.  — Arsenic  Pentasulfide,  260.  —  Sulfarsenites  and 
Sulfarsenates,  261.— Colloidal  Arsenic  Trisulfide,  261.  — Antimony,  Sb,  120.2. 
Occurrence  and  Preparation,  263. — Properties,  264.  —  Uses,  264.  —  Stibine, 

264.  —  Oxides  of  Antimony,  265.  —  Antimony  Hydroxide.    Antimonious  Acid, 

265.  —  Tartaric  Emetic,   266.  —  Antimonic  Acids,   267.  —  Chlorides  of  Anti- 
mony, 267.  —  Antimony  Trichloride,  267.  —  Antimony  Tetrachloride,  Hydro- 
tetrachloroantimonic  Acid,  267.  —  Antimony  Peutachloride,  267.  —  Metachlo- 
roantimonic  Acid,  268.  —  Antimony  Trisulfide,  268.  —  Antimony  Pentasulfide, 

268.  —  Sulfantimonites  and  Sulfantimonates,  268.  —Bismuth,  Bi,  208.     Occur- 
rence, Properties,  Uses,  268.  —  Oxides  of  Bismuth,  269.  —  Bismuth  Chloride, 

269.  —  Bismuth  Nitrate,  270.  —  Bismuth  Trisulfide,  270.  — Tables  of  Compounds 
of  the  Elements  of  the  Fifth  Group,  270.  —  Vanadium,  Columbium  Tantalum, 
271.  —  Exercises,  272. 

CHAPTER  XVI 
CARBON 

Carbon.  Occurrence,  273.  —  Diamonds,  274.  —  Graphite,  276.  —  Amorphous 
Carbon,  277. —Lampblack,  277.  — Wood  Charcoal,  277. —  Animal  Charcoal 
and  Bone  Black,  278.  —  Coke,  278.  —  Gas  Carbon.  Carbon  Electrodes,  279.  — 
Coal,  280.  —  Chemical  Properties  of  Carbon,  281. 


CHAPTER  XVII 

HYDROCARBOUS,    ILLUMINATING   AND    PRODUCER   GAS.    FLAME 

Marsh    Gas    or    Methane,    286.  —  Substitution,    287.  —  The    Davy    Safety 
Lamp,  287.  —  Homologues  of  Methane,  289.  —  Petroleum,  289.  —  Ethylene  or 


X  CONTENTS 

Ethene,  290.  —  Unsaturated  Compounds.  Ethylene  Chloride  and  Ethylene 
Bromide,  291.  —  Acetylene,  292.  —  Benzene,  294.  —  Illuminating  Gas,  295.  — 
Oil  Gas,  296.  —  Water  Gas,  296.  —  Producer  Gas,  297.  —  Blast-furnace  Gas, 
298.  —  Luminous  Flames,  299.  —  Bunsen  Burner,  300.  —  Explosion  Waves,  301. 
—  Temperature  of  Flames,  302.  —  Blowpipe,  303.  —  Reversed  Flames,  304.— 
Exercises,  305. 

CHAPTER  XVIII 


OXIDES   AND  SULPHIDES  OF  CARBON.    ASSIMILATION  AND 
RESPIRATION.    CYANIDES 

Carbon  Dioxide,  306.  —  Isothernials  of  Carbon  Dioxide,  307.  — Density  of 
Carbon  Dioxide,  308.  —  Aqueous  Solutions  of  Carbon  Dioxide,  Carbonic  Acid, 
309.  —  Carbonates  and  Bicarbonates.  Hard  Waters,  309.  —  Carbon  Monoxide, 
311.  — The  Cycle  of  Carbon  in  Nature,  312.  —  Respiration  Calorimeter,  313.— 
Carbon  Suboxide,  316.  —  Carbon  Oxychloride,  or  Phosgene  (Carbonyl  Chlo- 
ride), 316.  — Carbon  Bisulfide,  317.  — Sulfocarbonates,  317.  —  Sulfocarbonic 
Acid,  318.  —Carbon  Oxysulfide,  318.  —Cyanides,  319.  —  Hydrocyanic  Acid,  or 
Prussic  Acid,  319.  — Potassium  Cyanide,  319.  —  Complex  Cyanides,  319.— 
Potassium  Cyanate,  321.  —  Potassium  Thiocyanate,  321.  —  Cyanogen,  322. — 
Exercises,  322. 

CHAPTER  XIX 


ALCOHOLS,  ALDEHYDES,  KETONES,   ACIDS,   FATS,   CARBOHY- 
DRATES 

Structural  Formulas,  323.  —  l,  Valence,  323. — 2,  Radicals,  323.  —3,  Substi- 
tution, 324.  —  Alcohols,  324. —  Methyl  Alcohol,  324.  — Ethyl  Alcohol,  325.— 
Phenol  or  Carbolic  Acid,  326.  —  Glycerol,  326.  —  Aldehydes  and  Ketones,  327. 

—  Formaldehyde,  327.  — Benzaldehyde,  328.  — Acetone,  328.  — Acids,  328.— 
Formic  Acid,  329.  — Acetic  Acid,  329.  —  Oxalic  Acid,  329. —Lactic  Acid,  330. 

—  Tartaric  Acid,  330.  — Citric  Acid,  330.  —  Ammonium  Ferric  Citrate,  331.— 
Benzoic  Acid,  331.  — Palmitic,  Stearic  and  Oleic  Acids,  Fats,  331.  — Soaps, 
332.  —  Carbohydrates,  332.  —  Cane  Sugar,  or  Saccharose,  333.  —  Maltose,  334.  — 
Lactose,  or  Milk  Sugar,  334.  — Glucose,  334.— Fructose,  335. —  Starch,  335.— 
Dextrin,  336.  — Pectose,    Pectin,  337.  — Cellulose,   337.  — Paper,   337.  — Gun 
Cotton,  Celluloid,  Lacquers,  338. 


CHAPTER  XX 

AMINES,   DYES,   ALKALOIDS,    PROTEINS,    ENZYMES,    FOODS    AND 

NUTRITION 

Methyl  Amine,  339.  —  Aniline,  340.  — Dyes,  340.  — Alizarin,  341.  — Indigo, 
341.  —  Mordants,  342.  —  Alkaloids,  342.  —  Nicotine,  342.  —  Coniine,  342.  — 
Atropine,  343.  —  Cocaine,  343.  —  Morphine,  343.  —  Quinine,  343.  —  Strychnine, 
343. —Ptomaines,  343.  —  Proteins,  343.  —  Enzymes,  344.  —  Toxins,  Antitoxins, 
344.  —  Urea,  345.  —  Nutrition,  345. 


CONTENTS  Xi 

CHAPTER  XXI 

SILICON,  BORON,  GERMANIUM,  TIN,    LEAD,  TITANIUM,  ZIRCO- 
NIUM,  CERIUM,   THORIUM 

SILICON,  Si,  28.3 

Occurrence,  348.  — Preparation,  349.  —  Hydrogen  Silicide,  349.  — Silicon 
Carbide,  Carborundum,  349.  — Silicon  Fluoride,  350.  —  Fluosilicic  Acid,  350. 
Silicon  Tetrachloride,  351.— Silicon  Hexaiodide,  351.  — Silicon  Dioxide,  or 
Silica,  351.  — Artificial  Silicates,  352.  — Silicic  Acids,  353.  — Natural  Silicates, 
355. — Calculation  of  the  Formula  of  a  Mineral,  356. — Dialysis,  Semiperme- 
able  Membranes,  357.  — Osmotic  Pressure,  358.  — Germanium,  361. — Tin  and 
Lead,  361.  — Titanium,  362.  —  Zirconium,  363.  Cerium,  363.— Thorium,  364. 
—  Welsbach  Mantles,  364.  — Boron,  365.  — Preparation,  Properties,  365.— 
Boron,  Trioxide,  Borax  Beads,  365.  — Boric  Acid,  366.  — Other  Acids  of  Boron, 

366.  — Borax,  367. —Sodium  Perborate,  367.  — Other  Compounds  of  Boron, 

367.  — Exercises,  368. 


CHAPTER  XXII 

METALLIC    ELEMENTS.      DIFFERENCES    BETWEEN    METALS    AND 
NON-METALS.    PREPARATION  OF  COMPOUNDS 

Metals  and  Non-metals,  369.  —  Classification  of  the  Metals,  370.  — Melting 
Points  of  the  Elements,  372.  —  Preparation  of  Chemical  Compounds,  372.— 
Effect  of  Volatility,  374.  — Effect  of  Insolubility,  376.  — Effect  of  a  Common 
Ion.  Solubility  Product,  377.  —  Formation  of  Complex  Ions,  378. — Degree 
of  lonization,  379.  —  Effect  of  Degree  of  lonization,  Neutralization,  384. — 
Hydrolysis,  385.  — Illustration  of  the  Strength  of  Acids,  386.  — Use  of  Indi- 
cators, 387.  — Systematic  Study  of  the  Metals,  390.  — Metallurgy,  390.  — Ox- 
ides, 392.  — Hydroxides,  392.  —  Solubility  of  Salts,  393. 


CHAPTER  XXIII 

ALKALI  METALS:    LITHIUM,   SODIUM 

General  Properties  of  the  Alkali  Metals,  395.  — Lithium,  395.— Lithium 
Urate,  396.  — Atomic  Weight  of  Lithium,  Law  of  Diilong  and  Petit,  396.— 
The  Quantum  Theory,  398.  — Sodium,  398.— Metallurgy,  Properties,  399.— 
The  Alkali  Industry,  400.  — Sodium  Hydroxide,  401.  — Sodium  Oxide,  404. 
—  Sodium  Peroxide,  404.  — Sodium  Chloride,  404.  — Sodium  Sulfate,  Glauber's 
Salt,  406.  — Acid  Sodium  Sulfate  or  Sodium  Bisulf ate,  408. -Sodium  Sulfite, 
408. —Acid  Sodium  Sulfite,  or  Sodium  Bisulfite,  408. —Sodium  Hyposulfite, 
408.  — Sodium  Thiosulfate,  408.  — Sodium  Tetrathionate,  409. —  Sodium  Sul- 
fide,  409.  — Sodium  Hydrosulfide,  409.  — Sodium  Nitrate,  410. —Sodium  Ni- 
trite, 410.— Sodamide,  410. —  Sodium  Trinitride,  410.— Disodium  Phosphate, 
410.— Sodium  Carbonate  or  Sal  Soda  (Washing  Soda).  The  Leblanc  Soda 


xii  CONTENTS 

Process,  411.— Sodium  Bicarbonate  or  Baking  Soda,  The  Ammonia  Soda  Pro- 
cess, 412.  — Sodium  Silicate,  or  Soluble  Glass,  413.  — Sodium  Tetraborate,  or 
Borax,  413. 

CHAPTER  XXIV 

ALKALI  METALS:    POTASSIUM,  AMMONIUM,  RUBIDIUM,  CAESIUM. 
THE  SPECTROSCOPE 

Potassium.  Occurrence,  414.  —  Metallic  Potassium,  415.  —  Potassium  Oxide, 
415. — Potassium  Hydroxide,  415. — Potassium  Chloride,  416. —  Potassium 
Chlorate,  416.  — Potassium  Perchlorate,  416.— Potassium  Iodide,  417.  — Potas- 
sium Polyiodides,  417.  — Potassium  Sulfates,  417.  — Acid  Potassium  Sulfate, 
or  Potassium  Bisulfate,  417. —Potassium  Nitrate,  or  Saltpeter,  417.  — Gun- 
powder, 418.  — Potassium  Nitrite,  419.  — Potassium  Carbonate,  419.  —  Potas- 
sium Bicarbonate,  or  Saleratus,  420.  —  Potassium  Cyanide,  420.  —  Ammonium, 
420.  —  Ammonium  Hydroxide, 420.  —  Ammonium  Chloride,  421.  — Ammonium 
Sulfide,  421.  —  Ammonium  Hydrosulfide,  421.  —  Ammonium  Sulfate,  422.— 
Ammonium  Nitrate,  422.  —  Ammonium  Nitrite,  423.  —  Ammonium  Sodium 
Hydrogen  Phosphate,  423. —  Ammonium  Carbonate,  423.  —  Ammonium  Bi- 
carbonate, 423.  — Ammonium  Chloroplatinate,  423.  — Rubidium  and  Caesium, 
424.  — Spectrum  Analysis,  424. 


CHAPTER  XXV 

THE  ALTERNATE  METALS  OF  GROUP  I.    COPPER,  SILVER,  GOLD. 
PHOTOGRAPHY 

Copper.  Occurrence,  428.  —  Metallurgy,  428  —  Electrolytic  Refining  of  Cop- 
per, 429.  — Properties  of  Copper,  430.  —  Alloys  of  Copper,  431.  — Copper  Hy- 
droxide, 431.— Cupric  Oxide,  432.  — Cuprous  Oxide,  432.  — Cupric  Chloride, 
432.  — Cuprous  Chloride,  432. —Cuprous  Iodide,  433.— Cupric  Sulfide,  433.— 
Copper  Sulfate,  or  Blue  Vitriol,  433.  — Vitriols,  434. —Cupric  Nitrate,  434.— 
Ammoniocupric  Sulfate,  434.  —  Cuprous  Cyanide,  434.  —  Precipitation  of  Cop- 
per by  Iron,  Electromotive  Series,  435.  —  Faraday's  Law,  438.  —  Silver,  439. — 
Metallurgy,  439.  —  Pattison's  Process,  439.  —  Cupellation,  Assaying,  440. — 
Parke's  Process,  440.  —  Amalgamation  Process,  441.  — Other  Processes  for  the 
Recovery  of  Silver,  441.  — Properties  of  Silver,  Alloys,  442.  — Silver  Plating, 
442.  — Silver  Oxide,  442.  — Silver  Peroxide,  443.—  Silver  Nitrate,  444.  — Silver 
Nitrite,  444.  — Silver  Sulfate,  444.  — Silver  Chloride,  Silver  Bromide,  Silver 
Iodide,  444.  — Photography,  444.  — Gold,  445.  — Metallurgy,  446. —  Cyanide 
Process,  446.  —  Properties  of  Gold,  448.— Alloys  of  Gold,  448.  — Oxides  of 
Gold,  448. -Gold  Hydroxide,  448.  —  Chlorides  of  Gold,  450.  —Exercises,  450. 

CHAPTER  XXVI 

GROUP  II.    ALKALI-EARTH  METALS:   BERYLLIUM,  CALCIUM, 
STRONTIUM,  BARIUM,   RADIUM 

Beryllium,  451.  —Calcium.    Occurrence,  452.  —Preparation,  Properties,  452. 
-Calcium  Hydride,  452.  — Calcium  Oxide,  452.  — Dissociation  of   Calcium 


CONTENTS  xiii 

Carbonate  and  the  Phase  Rule,  453.  — Mortar,  454.  — Cement,  454.  — Cal. 
cium  Chloride,  455.— Chloride  of  Lime,  455.  — Calcium  Chlorate,  456.  — Cal- 
cium Flouride,  456.  — Calcium  Sulfide,  456.  — Acid  Calcium  Sulfite,  457.— 
Calcium  Sulfate,  Plaster  of  Paris,  457.  — Plaster  of  Paris  and  the  Phase  Rule, 
458.  — Calcium  Nitrate,  460.  — Calcium  Phosphates,  460.  —  Solubility  of  Cal- 
cium Phosphates,  461.— Calcium  Carbide,  462.  — Calcium  Cyanamide,  462.— 
Calcium  Carbonate,  463. —Hard  Waters,  463.— Determination  of  Free  and 
Combined  Carbonic  Acid  in  Natural  Waters,  464.  —  Calcium  Acetate,  465. — 
Calcium  Oxalate,  465.  —Calcium  Silicate,  466.  —  Glass,  466.  —  Strontium.  Oc- 
currence, 467.  — Strontium  Hydroxide,  468.  — Strontium  Nitrate,  468.  — Ba- 
rium. Occurrence,  468.  — Barium  Oxide,  468.  —  Barium  Peroxide,  469.— 
Barium  Hydroxide,  470.— Barium  Chloride,  470.  — Barium  Nitrate,  470.— 
Barium  Sulfide,  470.— Barium  Sulfate,  470.  — Flame  Colors  for  Calcium, 
Strontium  and  Barium,  471.  — Radium,  471. — Disintegration  of  Atoms,  472. 
—  Nature  of  the  Radiations  from  Radioactive  Substances,  473. — The  Life  of 
an  Element,  474.—  Other  Radioactive  Elements,  475.  — Chemical  Action  of  the 
Rays,  475.  —  Radiochemistry  in  Relation  to  Geology  and  Medicine,  475. — 
Exercises,  476. 


CHAPTER  XXVII 


ALTERNATE  METALS  OF  GROUP  II.     MAGNESIUM,   ZINC, 
CADMIUM  AND  MERCURY 

Magnesium,  478.  —  Preparation,  Properties,  478.  —  Magnesium  Oxide,  479. 

—  Magnesium  Hydroxide,  479. — Magnesium  Chloride,  480.  —  Magnesium  Am- 
monium Chloride,  480.  —  Magnesium  Sulfate,  480.  —  Magnesium  Sulfide,  480. 

—  Magnesium  Ammonium  Phosphate,  480.  —  Zinc.    Occurrence,  481.  — Metal- 
lurgy,  481.  — Uses.     Galvanized    Iron,  481.  —  Sherardized    Iron,  482.  — Zinc 
Oxide,  482.  — Zinc  Hydroxide,  483.  — Zinc  Chloride,  483. —  Zinc  Sulfate,  or 
White  Vitriol,  483.  — Zinc  Sulfide,  483.— Cadmium,  483.  — Cadmium  Hydrox- 
ide, 484.— Cadmium  Sulfate,  484.— Cadmium  Sulfide,  484.  —  Mercury,  Hg,  200.6. 
Occurrence.    Metallurgy,  484.  —  Properties  and  Uses,  485.  —  Amalgams,  486.  — 
Compounds  of  Mercury,  488.  — Mercurous  Oxide,  488.  —  Mercuric  Oxide,  488. 
Mercuric  Sulfide,   489.  —  Mercurous   Chloride,   or   Calamel,  489.  —  Mercuric 
Chloride,  or  Corrosive  Sublimates,  489.  —  Mercuric  Iodide,  490. — Mercurous 
Nitrate,  490.  —  Mercuric  Nitrate,  490.  —  Mercuric  Cyanide,  490.  —  Mercuric 
Fulminate,  491.  — lonization  of  Compounds  of  Cadmium  and  Mercury,  491.— 
Solubility  of  the  Sulfides  of  Group  II,  491.  —  Conduct  of  Solution  of  Mag- 
nesium, Zinc  and  Cadmium   Salts  towards  Ammonium  Hydroxide,   491.  — 
Ammono-mercuric  Compounds,  492.— Nessler's  Reagent,  492.— Exercises,  493. 


CHAPTER  XXVIII 

METALS  OF  GROUP  III.    ALUMINIUM  FAMILY.    RARE  EARTH 

METALS 

Aluminium,  494.  — Metallurgy,  495.  —  Properties  of  Aluminium,  497.— 
Alloys,  497.  —  Goldschmidt's  Thermite  Process,  497.  —  Aluminium  Chlo- 
ride, 498. 


xiv  CONTENTS 

— Aluminium  Fluoride,  499.  —  Aluminium  Hydroxide,  499. — Aluminium 
Oxide,  499.  — Aluminium  Sulfate,  500.  — Alums,  500.  — Brick,  Earthenware, 
Porcelain,  501.  — Ultramarine,  502.  — The  Rare  Earths,  502.  — Scandium,  503. 
Yttrium,  503. —  Lanthanum,  503. — Ytterbium,  504.  —  Praseodymium  and  Neo- 
dymium,  504.  —  Samarium,  505. — Europium,  Gadolinium,  Terbium,  505. — 
Holmium,  505.  — Dysprosium,  505.  — Erbium,  505.  —  Thulium,  506.  — Lute- 
cium, 506.  — Gallium,  506.  — Indium,  506.— Thallium,  507.  — Exercises,  507. 


CHAPTER  XXIX 

TIN  AND  LEAD 

Tin.  Occurrence,  Metallurgy,  508.  —  Uses  of  Tin,  Alloys,  Tin  Plate,  509.— 
Compounds  of  Tin,  510.  — Stannous  Oxide,  510.  — Stannous  Chloride,  510.— 
Stannous  Sulfide,  510.  —  Stannic  Oxide,  510.  —  Stannic  Acids,  511.  —  Stannic 
Acid,  511.  —  Metastannic  Acid,  512.  — Parastannic  Acid,  512.  — Stannic  Chlo- 
ride, 512.  — Stannic  Sulfide,  512.— Firep roofing  of  Cotton  Goods,  513.  — Lead. 
Occurrence,  Metallurgy,  513.  —  Properties  and  Uses  of  Lead,  Alloys,  514. — 
Oxides  of  Lead,  515. —  Lead  Monoxide,  or  Litharge,  515.  —  Storage  Batteries, 
516.— Lead  Sulfide,  518.— Lead  Chloride,  518.  — Lead  Tetrachloride,  518.— 
Lead  Sulfate,  519.— Lead  Nitrate,  519.  — Lead  Acetate,  or  Sugar  of  Lead,  519. 

—  Basic  Lead  Acetates,  519.— Lead  Carbonate,  519.  — Basic  Lead  Carbonate, 
or  White  Lead,  520. 

CHAPTER  XXX 

VANADIUM  AND  CHROMIUM  GROUPS 

Group  V.  Vanadium,  522.  —  Columbiurn  (or  Niobium),  523. — Tantalum, 
523. —Group  VI.  Chromium,  524.  —  Metallurgy,  Uses,  524.  —  Chromous 
Chloride,  525.  — Chromic  Oxide,  525.  — Chromic  Hydroxide,  525. —Chromic 
Chloride,  525.  —  Hydrates  of  Chromic  Chloride,  525. —  Potassium  Chromium 
Sulfate,  or  "Chrome  Alum,"  527.  —  Potassium  Chromate,  527. — Potassium 
Dichromate,  or  Pyrochromate,  527.  —  Lead  Chromate,  or  Chrome  Yellow,  527. 

—  Barium  Chromate,  528.  —  Chromium  Trioxide,  or  Chromic  Anhydride,  528. 
Chromyl  Chloride,  528.  —  Molybdenum,  528.  —  Molybdium  Trioxide,  or  Molyb- 
dic  Anhydride,  528.  —  Compounds  of  Molybdenum,  529.  —  Molybdic  Anhydride, 
529.  —Tungsten,  530.  —Compounds  of  Tungsten,  531.  —  Phosphotungstic  Acid, 
531.  — Uranium,  531. 


CHAPTER  XXXI 

MANGANESE 

Group  VII.  Manganese,  533.  — Occurrence,  Properties,  533.  —  Compounds 
of  Manganese,  534.  —  Manganous  Manganic  Oxide,  534.  —  Manganous  Hydrox- 
ide, 535. — Manganous  Chloride,  535.  —  Manganous  Sulfate,  535.  —  Manganous 
Sulfide,  535. — Manganese  Dioxide,  or  Black  Oxide  of  Manganese,  535.— 
Manganates,  536.  —  Permanganates,  537. — Potassium  Permanganate,  537. — 
Manganese  Heptoxide,  or  Permanganic  Anhydride,  538. 


CONTENTS  xv 


CHAPTER  XXXII 

IRON,  COBALT,  NICKEL 

Group  VIII.  Iron,  539.  —  Occurrence  of  Iron,  540. — Metallurgy  and  Iron, 
540.— Pig  Iron,  Cast  Iron,  543.— Wrought  Iron,  544.  —  Cementation  Steel, 
Cast  Steel,  545. —Bessemer  Steel,  547.  — Open  Hearth,  or  Siemens-Martin 
Process,  548.  — Alloy  Steels,  552.  — Compounds  of  Iron,  552.  — Potassium  Fer- 
rate, 553.  —  Ferrous  Chloride,  553.  —Ferrous  Hydroxide,  553.  —  Ferrous  Oxide, 

553.  —  Ferrous  Sulfate,  Green  Vitriol,  or  Copperas,  554.  —  Ferrous  Carbonate, 

554.  — Ferrous  Chloride  and  Nitric  Oxide,  554.  — Ferric  Chloride,  554.  — Ferric 
Hydroxide,  555.  —  Dialyzed  Iron,  555. — Ferric.  Oxide,  555. — Ferric  Sulfate, 
556.  — Magnetic  Oxide  of  Iron,  556.— Ferrous  Sulfide,  556. —Ferric  Sulfide, 
556.— Iron  Bisulfide,  or  Iron  Pyrites,  556.— Ferric  Thiocyanate,  556.— Co- 
balt,557. —  Compounds  of  Cobalt.    Oxides,  557.  —  Cobaltous  Hydroxide,  557. — 
Cobaltous  Chloride,  557.  — Cobalt  Sulfide,  558.  — Cobalt  Nitrate,  558.  — Cobalt 
Glass,  558.  — Potassium    Cobaltocyanide,    558.  — Potassium    Cobalticyanide, 
558.  —  Potassium  Cobaltinitrite,  558.  — Cobalt  Ammines,  559.  —Nickel,  559.  — 
Compounds  of  Nickel,  560.— Nickel  Dimethylglyoximine,  560.  — Nickel  Car- 
bonyl,  561. 


CHAPTER  XXXIII 

THE  PLATINUM  METALS 

Ruthenium,  563.  — Rhodium,  563.  —  Palladium,  563.  — Osmium,  564.— 
Iridium,  565.— Platinum,  565.  — Platinous  Chloride,  565.  — Chloroplatinic 
Acid,  565.— Platinic  Chloride,  566.— Platinum  Bisulfide,  566. 

INDEX       .  567 


A  TEXTBOOK  OF  CHEMISTRY 

CHAPTER  I 
INTRODUCTION 

The  Nature  of  Scientific  Knowledge.  The  phenomena  pre- 
sented to  our  senses  are  so  complex  and  varied  that  a  complete 
description  of  each  of  them  is  impossible.  It  is  the  purpose 
of  science  to  classify  these  phenomena  and  to  discuss  relation- 
ships between  them  which  are  frequently  repeated  and  separate 
them  from  those  relationships  which  are  not  repeated  and  which 
are  to  be  considered  as  more  or  less  accidental.  When  a 
given  relationship  between  two  or  more  phenomena  is  found 
always  to  exist,  we  conclude  that  the  relationship  is  necessary 
or  inherent  in  the  nature  of  things,  and  it  is  called  a  law.  The 
original  thought  conveyed  by  the  name  was,  doubtless,  that 
the  material  universe  acts  as  it  does  because  it  was  commanded 
to  do  so  by  some  higher  power,  but  the  word  has  come  to 
signify  simply  a  statement  of  some  constantly  recurring  rela- 
tionship between  phenomena. 

To  illustrate :  we  find  that  if  we  let  go  of  an  object  held  in 
the  hand  it  will  fall.  The  first  time  this  is  observed  it  is 
simply  a  fact  of  experience  or  observation.  But  a  further 
examination  of  the  relations  involved  leads  to  the  general 
statement  that  any  body  which  is  not  supported  will  fall. 
We  may,  perhaps,  call  this  a  law,  but  it  is  still  a  comparatively 
imperfect  statement  of  the  existing  relations.  A  further  study 
teaches  us  that  all  bodies  fall,  in  a  vacuum,  at  the  same  rate, 
irrespective  of  their  size  or  weight  and  that  the  velocity  of 
a  falling  body  varies  as  the  time  during  which  it  has  fallen. 
These  may  be  called  the  empirical  laws  of  gravitation,  that  is, 

1 


2  i  V    A  tSB00K  OF  CHEMISTRY 

the /laws  derived  m?m  ^experiment  and  observation.  For  the 
class  of  phenomena  to  which  they  apply  such  laws  have,  prac- 
tically, as  great  a  certainty  as  any  individual  fact  which  we  ob- 
serve. A  further  study  of  the  matter  and  especially  of  the 
motions  of  the  earth  and  planets  and  stars  has  led  to  the  more 
complete  generalization  that  all  bodies  attract  each  other  directly 
as  the  product  of  their  masses  and  inversely  as  the  squares  of 
their  distances.  This  is  Newton's  law  of  universal  gravitation. 
When  once  stated,  it  is  seen  that  the  empirical  laws  stated  above 
follow  from  it. 

A  natural  law,  when  discovered,  usually  requires  further  study 
and  development  in  two  directions.  It  calls,  in  the  first  place, 
for  a  careful  examination  of  the  logical  consequences  which  follow 
from  the  law,  and  it  furnishes  the  means  of  predicting  in  count- 
less cases  just  what  will  occur  in  given  conditions  where  the  law 
applies.  Thus  the  law  of  gravitation  is  assumed,  instinctively, 
every  time  that  we  move.  It  enables  us  to  predict,  accurately, 
the  motion  of  a  pendulum  or  projectile  and  to  calculate  the  exact 
relative  positions  of  the  sun,  moon  and  planets  for  a  hundred 
or  a  thousand  years  to  come.  The  law  lies  at  the  very  basis 
of  all  of  those  calculations  of  the  engineer  by  which  he  deter- 
mines the  necessary  strength  for  the  parts  of  a  bridge  or  a  truss. 
Many  other  illustrations  might  be  given  of  its  practical  impor- 
tance and  usefulness. 

In  the  second  place,  a  natural  law  suggests  the  need  of  some 
additional  explanation.  An  inquiry  in  this  direction  may  lead 
to  some  more  fundamental  and  general  law,  as  when  the  laws  of 
falling  bodies  were  found  to  be  only  special  cases  of  the  universal 
law  of  gravitation,  or  it  may  lead  to  the  confines  of  our  present 
knowledge  at  a  point  where  no  further  progress  can  be  made 
without  the  use  of  speculation  or  hypothesis.  Following  our 
illustration,  it  appears  to  most  minds  almost  or  quite  incon- 
ceivable that  one  body  should  act  upon  another  at  a  distance 
without  some  medium,  and  there  have  been  many  speculations 
with  regard  to  some  medium  which  may  be  the  cause  of  gravita- 
tion. Several  hypotheses  with  regard  to  the  mechanism  of  the 


INTRODUCTION  3 

action  of  such  a  medium  have  been  proposed,  but  these  specula- 
tions have  not  met  with  any  notable  success. 

With  regard  to  such  speculation  there  are,  at  present,  among 
scientific  men,  and  notably  among  chemists,  two  somewhat  dis- 
tinct schools  or  classes.  One  of  these  takes  the  ground  that  the 
number  of  possible  explanations  of  those  parts  of  the  universe 
which  lie  beyond  our  knowledge  is  so  great  that  it  is  hopeless 
to  attempt  to  find  the  true  explanation.  While  admitting  the 
value  of  hypotheses  in  stimulating  and  directing  research,  this 
school  claims  that  such  hypotheses  can  never  give  us  any  real 
knowledge  of  matters  which  are  beyond  the  cognizance  of  our 
senses,  and  that  all  genuine  scientific  advance  consists  in  giving 
a  fuller  description  of  things  about  which  we  can  gain  direct, 
positive  knowledge.  The  other  school  points  out  that,  while 
there  are  many  things  in  the  universe  which  must  always  remain 
beyond  the  possibility  of  direct  knowledge,  we  can  accumulate 
so  much  evidence  with  regard  to  these  that  it  may  be  possible, 
ultimately,  to  give  to  our  theories  with  regard  to  them  a  high 
degree  of  probability.  The  danger  of  the  first  attitude  of  mind 
is  that  the  investigator  will  be  content  with  a  full  description  of 
phenomena  and  will  fail  to  discover  relations  which  can  be  under- 
stood only  by  a  knowledge  of  matters  about  which  we  can  secure 
only  indirect  evidence.  The  danger  of  the  other  point  of  view 
is  that  it  may  lead  one  to  overestimate  the  amount  of  knowledge 
which  has  been  acquired  about  unseen  things  and  to  spend  time 
in  useless  speculations  which  would  be  better  spent  in  acquiring 
new  facts. 

Whichever  view  is  accepted,  the  science  of  our  time  includes 
a  knowledge  of  a  very  great  number  of  facts,  of  the  natural  laws 
which  express  the  relations  between  these'  facts  and  of  the 
theories  which  are  our  best  present  explanation  of  the  laws. 

Subdivisions  of  Science.  There  is,  properly  speaking,  only 
one  science,  which  includes  all  classified,  systematic  knowledge. 
The  amount  of  such  knowledge  has  become  so  great,  however, 
that  it  is  customary  to  subdivide  it  into  a  number  of  parts,  each 
of  which  is  called  a  science.  It  should  always  be  remembered 


A  TEXTBOOK  OF  CHEMISTRY 


that  the  boundaries  between  various  divisions  are  more  or  less 
arbitrary  and  that  very  many  facts  belong  about  equally  to  two 
or  more  of  the  sciences.  It  is  also  very  important  to  understand 
that  no  one  can  make  much  progress  in  any  one  science  without 
considerable  knowledge  of  several  others. 

The  more  important  subdivisions  of  science  are  the  following : 
abstract  sciences,  which  deal,  primarily,  with  forms  of  abstract 
reasoning  —  mathematics  and  logic;  physical  sciences,  dealing 
with  the  phenomena  of  matter  and  energy  apart  from  life  — 
physics,  chemistry,  astronomy,  mineralogy,  geology;  biological 
sciences,  dealing  with  the  phenomena  of  living  bodies  —  bacteri- 
ology, botany,  zoology,  paleontology;  psychological  sciences, 
dealing  with  the  phenomena  of  mind  and  of  society — psychology, 
language,  history,  social  science,  political  economy,  ethics. 

Mathematics 
Abstract  sciences 

Logic 

Physics 

Chemistry 

Astronomy 

Mineralogy 

.  Geology 

Bacteriology 

Botany 

Zoology 

Paleontology 

Psychology 

Language 

History 

Social  science 

Political  economy 

.Ethics 

Physical  Sciences.  The  two  fundamental  physical  sciences 
are  physics  and  chemistry.  The  other  three  mentioned,  astron- 
omy, mineralogy  and  geology,  are  concerned  with  the  applica- 
tion of  the  laws  of  physics  and  chemistry  in  studying  particular 
bodies  or  substances  and  so  are  more  special  in  their  nature. 


Physical  sciences 


Biological  sciences 


Psychological  sciences 


INTRODUCTION  5 

Roughly  speaking,  physics  treats  of  energy  and  chemistry  of 
matter.  Thus  chemistry  tells  us  of  the  properties  and  composi- 
tion of  substances,  as  of  water,  of  iron  or  of  sulfur,  of  the  action 
of  substances  on  each  other  and  of  the  changes  in  composition 
which  they  undergo  in  a  great  variety  of  circumstances.  Phys- 
ics, on  the  other  hand,  deals  with  the  varieties  of  energy,  as 
mechanical  energy,  sound,  heat,  light,  electricity,  and  with  the 
transformations  of  each  of  these  into  other  forms. 

As  we  can  have  no  knowledge  of  matter  except  through  the 
energy  which  it  possesses  and  the  effect  of  that  energy  on  our 
senses  in  one  way  or  another,  and  since  the  changes  in  energy 
which  result  when  substances  act  on  each  other  are  often  of  great 
importance,  the  chemist  can  make  little  progress  in  his  study 
without  a  considerable  knowledge  of  physics.  And  as  we  have 
no  knowledge  of  energy  apart  from  matter,  the  physicist  finds 
some  knowledge  of  chemistry  desirable.  This  interrelation  be- 
tween the  sciences  is  so  close,  also,  that  there  is  a  large  domain 
which  is  common  to  both  and  of  which  it  is  scarcely  worth  while 
to  ask  whether  it  belongs  to  chemistry  or  to  physics. 

Matter  and  Energy.  The  two  most  fundamental  concepts  of 
physical  science  are  matter  and  energy.  Matter  may  be  defined 
as  anything  which  has  mass,  or,  in  its  relation  to  the  earth  or  to 
other  bodies,  weight.1  Putting  the  same  thought  in  quite  different 
words,  matter  is  anything  which  requires  energy  to  set  it  in  mo- 

1  This  definition  of  matter,  which  is  based,  of  course,  on  Newton's 
conception  of  inertia  —  that  no  body  can  move,  if  at  rest,  or  change 
the  direction  or  velocity  of  its  motion,  if  moving,  unless  it  is  acted 
on  by  some  external  force  —  is  not  entirely  satisfactory.  It  has 
been  shown  that  the  mass  of  a  body  is  changed  by  a  change  in  its 
velocity,  though  the  change  is  inappreciable  until  the  velocity 
approaches  that  of  light.  (D.  F.  Comstock,  J.  Amer.  Chem.  Soc.,  30, 
683).  In  such  a  case  it  is  usually  better  to  retain  the  older,  simple 
definition,  frankly  recognizing  that  it  is  imperfect.  It  is  character- 
istic of  a  scholastic  rather  than  a  scientific  attitude  of  mind  to  be 
much  troubled  because  a  definition  is  imperfect  or  incomplete.  The 
scientific  worker  sees  that  our  knowledge  is  incomplete  in  every 
direction  and  is  constantly  developing.  The  definitions  are  simply 
a  means  of  conveying  this  incomplete  knowledge  to  others  and  of  at- 
tempting to  discover  the  most  fundamental  conceptions  of  science. 
We  succeed  best  in  both  directions  by  keeping  these  definitions 
simple. 


6  A   TEXTBOOK  OF  CHEMISTRY 

tion  or  to  change  its  rate  of  motion.  Energy  is  usually  defined, 
on  the  basis  of  etymology,  as  anything  which  can  do  work.  A 
more  satisfactory  definition  is  that  energy  is  anything  which  may 
set  matter  in  motion  or  change  its  rate  of  motion.  The  most 
important  forms  of  energy  are  mechanical  energy,  sound,  heat, 
light,  electrical  energy  and  chemical  energy. 

Conservation  of  Matter  and  Energy.  A  superficial  observation 
of  many  phenomena  in  nature  appears  to  show  that  under  some 
conditions,  especially  in  burning,  matter  is  destroyed,  and  that 
under  other  conditions,  as  in  the  rusting  of  iron,  matter  in- 
creases in  weight.  It  is  comparatively  easy  to  show,  however, 
that  a  part  of  the  air,  which  has  weight,  takes  part  in  these  pro- 
cesses and  that  while  a  candle,  for  instance,  seems  to  be  destroyed 
in  burning,  water  and  carbon  dioxide  are  formed  by  its  combus- 
tion ;  and  if  these  are  absorbed  by  soda  lime  and  weighed,  the  sum 
of  their  weights  is  very  considerably  greater  than  the  weight  of 
the  candle  which  has  been  burned.  Still  more  careful  experi- 
ments will  show  that  the  products  of  combustion  weigh  exactly 
the  same  as  the  weight  of  the  candle  and  the  weight  of  the  por- 
tion of  the  air  (oxygen)  with  which  it  has  combined. 

The  question  whether  there  is  any  change  in  the  weight  of 
matter  during  a  chemical  reaction  is  so  fundamental  that  one 
chemist  (Landolt)  has  considered  it  worth  while  to  give  ten  years 
of  most  careful  and  painstaking  work  to  its  study.  His  con- 
clusion is  that  in  the  cases  which  he  studied  no  change  so  great 
as  the  one  millionth  part  of  the  weight  of  the  substances  which 
reacted  with  each  other  occurred.  We  say,  therefore,  that  no 
method  is  known  by  which  we  can  create,  or  destroy,  matter. 
This  is  known  as  the  law  of  conservation,  or  indestructibility  of 
matter. 

If  we  place  a  wheel  with  vanes  in  a  can  of  water  and  wind 
around  its  axle  a  cord  tied  to  a  weight  in  such  a  manner  that  as 
the  weight  falls  the  wheel  will  revolve  and  stir  the  water,  we  shall 
find  that  the  temperature  of  the  water  will  rise.  If  the  experi- 
ment is  carefully  performed,  it  will  be  found  that  a  weight  of 
one  kilogram  falling  427  meters  will  raise  the  temperature  of  a 


INTRODUCTION  7 

kilogram  of  water  one  degree.  On  the  other  hand,  if  the  steam 
from  a  boiler  is  caused  to  drive  the  piston  of  a  steam  engine 
which  is  pumping  water  or  doing  other  work,  it  is  found  that  a 
part  of  the  heat  of  the  steam  disappears  and  that  the  heat  which 
can  no  longer  be  found  in  the  exhaust  steam  or  anywhere  about 
the  engine  corresponds  accurately  to  the  amount  of  work  which 
the  engine  performs,  and  that  the  ratio  is  exactly  the  same  as 
that  found  between  the  falling  weight  and  the  rise  in  tempera- 
ture of  water  stirred  by  the  paddle  wheel.  The  energy  of  the 
engine  may  be  used  to  drive  a  dynamo  which  will  furnish  an  elec- 
tric current ;  the  electric  current  may  be  used  to  decompose  a 
chemical  compound,  giving  substances  which  contain  more  chem- 
ical energy  than  the  compound ;  and  these  substances,  in  turn, 
may  be  recombined,  giving  out  heat  in  the  process.  Each  form 
of  energy  which  we  know  may  be  transformed  into  some  other, 
and  there  is  always  an  exact  relation  between  the  quantity  of 
energy  of  one  kind  which  disappears  and  the  quantity  of  other 
kinds  of  energy  which  takes  its  place.  This  is  the  law  of  the 
conservation  of  energy.  It  might  be  called  the  law  of  the  inde- 
structibility of  energy. 

Pure  Substances  and  Mixtures.  As  has  been  stated,  chemistry 
deals,  primarily,  with  the  properties  and  composition  of  sub- 
stances.1 If  we  examine  certain  substances,  such  as  pure  water 
or  gold,  we  find  that  they  are  alike  throughout  their  whole  mass, 
or  that  they  are  homogeneous.  We  find  that  every  sample  of 
such  a  substance  which  we  examine  melts  or  freezes  at  exactly 
the  same  temperature ;  and  that  if  it  boils  without  decomposition, 
it  will  always  boil  at  the  same  temperature  under  atmospheric 
pressure.  We  find,  too,  that  if  the  substance  is  a  liquid  or  gas, 
the  density  or  specific  gravity  is  always  the  same  under  the  same 
condition  of  temperature  and  pressure.  In  spite  of  the  large 

1  The  distinction  between  the  words  body  and  substance  should 
be  carefully  observed.  Body  always  refers  to  some  definite,  con- 
crete thing,  as  a  heavenly  body,  speaking  of  the  sun  or  a  star, 
a  body  of  ore,  etc.  Substance,  on  the  other  hand,  refers  to  some  par- 
ticular kind  of  matter,  as  water  or  gold.  A  given  piece  of  gold 
might  be  called  a  body,  but  gold,  in  general,  is  a  substance. 


8  A  TEXTBOOK  OF  CHEMISTRY 

number  of  substances  which  are  known  to  exist  (more  than  one 
hundred  thousand)  it  is  possible  to  identify  many  of  these  with 
practical  certainty  by  the  examination  of  a  comparatively  small 
number  of  their  properties.  For  instance  no  other  substance  has 
the  same  freezing  point,  boiling  point  and  density  as  water. 

Contrasted  with  pure  substances,  as  water  or  gold,  most  sub- 
stances which  we  meet  in  daily  experiences  are  mixtures.  For 
example,  if  we  take  a  cereal,  as  wheat,  and  powder  it,  as  is  done 
in  the  milling  process,  a  portion  will  pass  through  fine  bolting 
cloth,  while  another  portion,  the  bran,  will  not.  If  the  portion 
which  passes  the  bolting  cloth  is  warmed  gently,  it  loses  weight, 
and  it  can  readily  be  shown  that  the  loss  is  almost  wholly  due 
to  the  escape  of  water,  which  may  be  condensed  and  identified 
by  its  freezing  point  and  boiling  point.  From  the  dry  flour  ether 
will  dissolve  an  oil  or  fat  which  will  be  left  behind  on  evaporating 
the  ether.  If  the  portion  which  remains  is  kneaded  between  the 
fingers  in  a  stream  of  running  water,  a  fine  white  powder,  consist- 
ing mainly  of  starch,  will  be  washed  away,  while  a  residue,  called 
gluten,  which  consists  largely  of  proteins,  will  remain.  The 
processes  described  show  clearly  that  the  cereal  is  a  very  complex 
mixture  of  many  different  substances,  but  of  the  substances  sepa- 
rated only  the  water  and  starch  can  be  considered  as  even  approx- 
imately pure  substances.  A  more  careful  examination  of  the  bran 
or  oil  or  gluten  will  show  that  each  of  these  is  still  a  mixture. 

Preparation  of  Pure  Substances.  A  large  part  of  the  work 
which  must  be  done  in  the  study  of  chemistry  consists  in  the 
separation  and  characterization  of  pure  substances.  The  most 
common  means  used  for  this  purpose  are  treatment  of  mixtures 
with  solvents,  crystallization  and  distillation.  Thus  if  we  have 
a  mixture  of  sugar  and  sand,  the  sugar  may  be  easily  separated 
by  dissolving  it  in  water  and  pouring  off  or  filtering  the  solution 
from  the  sand.  From  a  brine  which  contains  other  substances 
in  solution  along  with  salt,  the  salt  may  be  obtained  nearly  pure 
by  evaporating  it  till  the  salt  separates  in  crystals.  For  some 
reason  particles  of  the  same  kind  separate  from  a  solution  on 
evaporation  or,  frequently,  on  cooling  a  hot  solution,  in  definite, 


INTRODUCTION 


geometrical  forms  called  crystals,  and  when  they  separate  in 
this  manner  they  usually  exclude  other  substances  which  may 
be  present.  By  repeated  distillation  of  a  mixture  of  alcohol  and 
water,  collecting  the  lower  boiling  portions  by  themselves  each 
time,  nearly  pure  alcohol  can  be  separated  from  water.  When  a 
volatile  substance  contains  a  nonvolatile  one  in  solution,  the 
separation  by  distillation  is  much  easier  and  , — , 

more  complete. 

Elements  and  Compounds.  If  the  red 
oxide  of  mercury  is  heated  in  a  small  tube, 
metallic  mercury  will  distill  away,  while  a 
glowing  splinter  held  at  the  mouth  of  the 
tube  will  burst  into  flame.  The  heat  causes 
the  decomposition  of  the  oxide  of  mercury 
into  mercury  and  a  gas  which  supports  com- 
bustion better  than  air,  and  which  is  called 
oxygen.  An  electric  current  passed  between 
two  strips  of  platinum  immersed  in  a  solution 
of  sulfuric  acid  in  water  in  the  apparatus  -^ 
shown  in  Fig.  1  will  cause  the  separation  of 
two  gases,  oxygen  and  hydrogen.  As  it  can 
be  shown  that  the  amount  of  sulfuric  acid 
remains  unchanged,  it  is  evident  that  the 
gases  are  formed  by  the  decomposition  of 
the  water;  and  this  view  can  be  confirmed 
by  burning  the  mixture  of  oxygen  and  hydrogen  and  regenerat- 
ing the  water.  While  oxide  of  mercury  can  be  decomposed 
into  mercury  and  oxygen,  and  water  may  be  decomposed  into 
oxygen  and  hydrogen,  no  one  has  thus  far  succeeded  in  decom- 
posing mercury  or  oxygen  or  hydrogen.  Substances  like  these, 
which  it  has  not  been  found  possible  to  decompose,  are  called 
elements.1  Substances  which  can  be  separated  into  two  or 

1  This  definition  is  not  wholly  satisfactory,  since  it  has  been 
found  that  radium,  which  has  all  of  the  other  properties  of  an  ele- 
ment, decomposes  spontaneously  into  helium  and  a  whole  series 
of  other  elements.  It  seems  best,  however,  to  retain  the  simple 
definition,  but  also  it  is  best  to  consider  radium  as  an  element. 


J 


Fig.  1 


10 


A  TEXTBOOK  OF  CHEMISTRY 


more  parts,  neither  of  which  can  be  converted  into  the  other,  or 
which  can  be  prepared  by  the  union  of  two  or  more  elements, 
are  called  compounds.  Elements  sometimes  exist  in  two  or 
three  different  forms,  but  these  may  always  be  converted  each 
into  the  other. 

Only  about  eighty  elements  have  been  positively  identified. 
The  names  of  these,  together  with  their  symbols  and  atomic 
weights,  are  given  in  the  following  table : 

ATOMIC 
WEIGHT 

27.1 
120.2 

39.88 

74.96 
137.37 
208.0 

11.0 

79.92 
112.40 
132.81 

40.07 

12.00 
140.25 

35.46 

52.0 

58.97 

93.5 

63.57 
162.5 
167.7 
152.0 

19.0 
157.3 

69.9 

72.5 
9.1 
197.2 

3.99 
163.5 

1.008 
114.8 
126.92 
193.1 

1Also  called  niobium,  Nb.  2  Often  given  as  beryllium,  Be. 


SYMBOL 

Aluminium 

.   Al 

Antimony    . 

.   Sb 

Argon      .     . 

.   A 

Arsenic   . 

.   As 

Barium   .     . 

.   Ba 

Bismuth 

.   Bi 

Boron     .     . 

.    B 

Bromine 

.    Br 

Cadmium    . 

.    Cd 

Caesium 

.   Cs 

Calcium 

.   Ca 

Carbon   .     . 

.   C 

Cerium   .     . 

.   Ce 

Chlorine 

.   Cl 

Chromium  . 

.   Cr 

Cobalt     .     . 

.   Co 

Columbium  1 

.   Cb 

Copper    .     . 

.   Cu 

Dysprosium 

•   Dy 

Erbium   .     . 

.   Er 

Europium    . 

.   Eu 

Fluorine 

.   F 

Gadolinium 

.   Gd 

Gallium 

.   Ga 

Germanium 

.   Ge 

Glucinum2  . 

.    Gl 

Gold  .     .     . 

.  Au 

Helium   .     . 

.   He 

Holmium     . 

.   Ho 

Hydrogen    . 

.   H 

Indium   .     . 

.   In 

Iodine     .     . 

.   I 

Iridium   .     . 

.   Ir 

ATOMIC 

SYMBOL 

WEIGHT 

Iron  .     ... 

Fe 

55.84 

Krypton      .     . 

Kr 

82.92 

Lanthanum 

La 

139.0 

Lead        .     .     . 

Pb 

207.10 

Lithium       .     . 

Li 

6.94 

Lutecium     .     . 

Lu 

174.0 

Magnesium 

Mg 

24.32 

Manganese 

Mn 

54.93 

Mercury      ,     . 

Hg 

200.6 

Molybdenum    . 

Mo 

96.0 

Neodymium 

Nd 

144.3 

Neon       .     .     . 

Ne 

20.2 

Nickel     .     .     . 

Ni 

58.68 

Niton      .     .-    . 

Nt 

222.4 

Nitrogen      .   .  . 

N 

14.01 

Osmium       .     . 

Os 

190.9 

Oxygen  .     .     . 

0 

16.00 

Palladium    .     . 

Pd 

106.7 

Phosphorus 

P 

31.04 

Platinum     .     . 

Pt 

195.2 

Potassium 

K 

39.10 

Praseodymium 

Pr 

140.6 

Radium  .     .     . 

Ra 

226.4 

Rhodium 

Rh 

102.9 

Rubidium    .     . 

Rb 

85.45 

Ruthenium 

Ru 

101.7 

Samarium   . 

Sa 

150.4 

Scandium    .     . 

Sc 

44.1 

Selenium      .     . 

Se 

79.2 

Silicon     .     .     . 

Si 

28.3 

Silver      .     .     . 

Ag 

107.88 

Sodium        .     . 

Na 

23.00 

Strontium    .     . 

Sr 

87.63 

INTRODUCTION 


11 


ATOMIC 

ATOMIC 

SYMBOL      WEIGHT 

SYMBOL 

WEIGHT 

Sulfur      .     . 

.  S 

32.07 

Uranium 

.  U 

238.5 

Tantalum     . 

.  Ta 

181.5 

Vanadium    . 

.  V 

51.0 

Tellurium     . 

.  Te 

127.5 

Xenon     .     . 

.  Xe 

130.2 

Terbium 

.  Tb 

159.2 

Ytterbium 

Thallium      . 

.  Tl 

204.0 

(Neoytter- 

Thorium 

.  Th 

232.4 

bium)   . 

.  Yb 

172.0 

Thulium      . 

.  Tm 

168.5 

Yttrium 

.  Y 

89.0 

Tin     .     .     . 

.  Sn 

119.0 

Zinc    .     .     . 

.  Zn 

65.37 

Titanium     . 

.  Ti 

48.1 

Zirconium    . 

.  Zr 

90.6 

Tungsten     . 

.  W 

184.0 

The  symbols  are  either  the  first  letter  or  the  first  letter  together 
with  some  other  characteristic  letter  of  the  name  of  the  element. 
With  few  exceptions  symbols  are  derived  from  the  English  names 
and  the  symbols  readily  suggest  the  names.  The  exceptions  are : 


Antimony,  Sb,  Stibium 


Potassium,  K,    Kalium 


Gold, 
Iron, 
Lead, 
Mercury, 


Au,  Aurum 
Fe,  Ferrum 
Pb,  Plumbum 
Hg,  Hydrargyrum 


Silver,          Ag,  Argentum 
Sodium,       Na,  Natrium 
Tin,  Sn,  Stannum 

Tungsten,    W,    Wolfram 
For  all  of  these  except  the  last  the  symbols  are  derived  from  the 
Latin  names. 

The  elements  vary  greatly  in  their  relative  abundance.  Of  that 
portion  of  the  earth  which  we  are  able  to  examine  it  is  estimated 
that  oxygen  forms  nearly  one  half  of  the  total  weight  and  sili- 
con one  fourth.  The  percentage  amounts  of  the  twelve  most 
common  elements  in  the  surface  of  the  earth  to  a  depth  of  ten 
miles,  including  the  ocean  and  the  atmosphere,  are  estimated  as 

follows : l 

PER  CENT  PER  CENT 

Oxygen,  49.78  Potassium,         2.28 

Silicon,  26.08  Magnesium,       2.24 

Aluminium,       7.34  Hydrogen,         0.95    . 

Iron,  4.11  Titanium,          0.37 

Calcium,  3.19  Chlorine,  0.21 

Sodium,  2.33  Carbon,  0.19 

99.07 
1  F.  W.  Clarke,  Data  of  Geochemistry,  p.  32. 


12  A  TEXTBOOK  OF  CHEMISTRY 

Some  elements  which  form  only  a  very  small  part  of  the  whole 
are  very  important,  especially  nitrogen,  phosphorus  and  several 
of  the  metals  which  are  not  included  in  the  above  table. 

How  Pure  Substances  are  distinguished  from  Mixtures. 
Law  of  Constant  Proportion.  A  very  large  part  of  our  knowledge 
of  chemistry  depends  on  the  preparation  of  pure  substances  and 
on  the  determination  of  the  properties  and  composition  of  these. 
It  is,  therefore,  important  to  understand  how  we  may  distinguish 
between  pure  substances  and  mixtures.  The  first  characteristic 
of  a  pure  substance  is  that  it  must  be  homogeneous  so  long  as 
it  exists  in  one  state  of  aggregation,  that  is,  so  long  as  it  is  all 
solid,  all  liquid  or  all  gaseous.  Second,  it  must  have  a  constant 
melting  point  and  boiling  point,  if  it  melts  and  boils  without  de- 
composition, and  the  specific  gravity  or  density  and  other  physi- 
cal properties  must  be  invariable  under  the  same  conditions.1 
Third,  a  pure  substance  must  always  show  the  same  conduct 
toward  any  other  substance  which  may  dissolve  it  or  act  upon 
it  chemically,  provided  that  the  conditions  are  the  same. 

A  very  careful  examination  of  a  large  number  of  substances 
which  have  the  characteristics  just  given  in  the  highest  degree 
has  demonstrated  that  such  substances  are  absolutely  constant 
in  composition.  This  is  the  law  of  constant  proportion  and  may 
be  stated  thus :  A  pure  substance  always  contains  the  same  ele- 
ments in  the  same  proportion  by  weight.  Thus  pure  water  always 
contains  hydrogen  and  oxygen  in  the  proportion  of  1  to  7.94 
parts  by  weight.  This  law  has  been  tested  by  a  large  amount 
of  most  careful  and  painstaking  work,  and  the  more  careful  the 
work  has  been  the  more  accurately  has  the  law  been  found  true, 
so  that  we  may  consider  it  as  one  of  the  most  absolutely  perfect 
laws  of  nature.  Since  a  very  large  number  of  substances  which 
fulfill  the  first  three  requirements  of  a  pure  substance  are  in- 
variable in  composition,  this  constancy  of  composition  is  con- 
sidered as  a  fourth  characteristic  of  a  pure  substance.  It  is  a 

1  The  density  of  some  solids  and  especially  of  metals  may  vary 
slightly  according  to  the  treatment  to  which  they  have  been  sub- 
jected. 


INTRODUCTION  13 

characteristic  of  very  great  importance  and  one  which  is  fre- 
quently used  to  determine  whether  a  given  substance  is  pure  or 
not. 

Inductive  Reasoning.  It  may  seem  at  first  that  the  use  of 
constancy  of  composition  as  a  means  of  determining  whether  a 
substance  is  pure  or  not  is  due  to  reasoning  in  a  circle,  or,  as  it  is 
commonly  called,  is  "  begging  the  question."  We  say  first  that 
a  pure  substance  has  a  constant  composition  and  then  that 
because  a  substance  has  a  constant  composition  it  is  pure.  The 
criticism  would  be  justified  if  constancy  of  composition  were 
the  only  characteristic  applied  to  decide  whether  a  substance  is 
pure  or  not.  But  the  first  three  characteristics  mentioned  above 
are  the  ones  which  will  appeal  to  any  one  as  being  dictated  by 
common  sense.  When  we  find  that  a  very  great  number  of 
substances  having  these  characteristics  are  also  constant  in  com- 
position, we  come  to  the  conclusion  that  there  is  some  inherent, 
necessary  connection  between  this  fourth  characteristic  and  the 
other  three,  and  that  when  a  given  substance  does  not  have  this 
characteristic  it  probably  lacks  some  of  the  other  three  as  well. 
Such  a  conclusion  is  said  to  be  reached  by  inductive  reasoning. 
The  truth  of  such  a  conclusion  can  never  be  absolutely  proved 
any  more  than  we  can  prove  that  the  sun  will  rise  to-morrow 
morning.  But  we1  may  reach  practical  certainty  by  means  of 
such  conclusions  and  may  properly  use  them  as  the  basis  for 
further  reasoning. 

Law  of  Combining  Weights.  If  we  select  a  series  of  com- 
pounds in  such  a  manner  that  each  compound  has  an  element 
contained  in  the  preceding  and  another  contained  in  the  follow- 
ing compound,  it  will  be  found  that  whenever  the  same  element 
recurs  the  proportion  of  the  element  which  combines  with  other 
elements  will  always  be  the  same  or  some  exact  multiple  or 
submultiple  of  the  first  proportion.  This  will  be  more  clear 
from  the  following  series  of  compounds  *• l 

1  Here  and  elsewhere  whole  numbers  are  used  for  greater  sim- 
plicity. The  exact  values  will  be  found  in  the  table  of  atomic 
weights,  p.  10. 


14 

A  TEXTBOOK 

OF  CHEMISTRY 

Water 

Cuprous 

Cupric             Hydrogen      Hydrochloric 

Oxide 

Sulfide 

Sulfide 

Acid 

H:O 

0:Cu 

Cu 

:S 

S:H 

H:C1 

1:8 

8  :  63.6 

63.6 

:32 

32:2 

2:71 

Ferrous 

Ferrous 

Sulfur 

Sodium 

Sodium 

Sodium 

Chloride 

Oxide           Dioxide 

Sulfide 

Chloride 

Chlorate 

Cl:Fe 

Fe:O 

0:S 

S:Na 

Na:Cl 

Na  :  Cl  : 

O 

71:56 

56:16 

16:16 

16:23 

23  :  35.5 

23  :  35.5  :  48 

In  this  series  of  compounds  hydrogen  has  been  chosen  as  the 
starting  point  and  has  been  given  a  value  of  1.  If  oxygen  had 
been  chosen  and  had  been  given  a  value  of  100,  as  was  at  one 
time  proposed,  the  other  numbers  would  all  be  different  but 
exactly  the  same  ratios  between  the  different  numbers  for  the 
same  element  would  be  found  throughout  the  series. 

It  is  seen  that  the  values  for  hydrogen  are  1  and  2,  for  oxygen 
8,  16,  32,  and  48,  for  sulfur  32  and  16,  for  chlorine  71  and  35.5, 
the  larger  numbers  for  each  element  being  in  every  case 
exact  multiples  of  the  smallest  number  for  the  element. 
This  table  might  be  extended  to  include  all  pure  substances  which 
have  been  analyzed.  The  law  of  combining  weights  stated 
above  may  be  expressed  more  briefly  as  follows :  A  number 
may  be  selected  for  each  element  which  represents  the  proportion 
of  the  element  which  enters  into  combination  with  other  elements. 

The  Atomic  Theory.  The  laws  of  constant  proportion  and 
of  combining  weights  find  a  very  satisfactory  explanation  in  the 
atomic  theory,  which  was  proposed  by  Dalton  at  the  beginning 
of  the  nineteenth  century.  According  to  this  theory  the  chemi- 
cal elements  are  composed  of  very  small  particles  or  atoms,  the 
atoms  of  the  same  element  being  all  alike  in  properties  and  in 
weight,  while  the  atoms  of  different  elements  are  different.  If 
we  suppose  further  that  compounds  are  always  formed  by  the 
union  of  atoms  of  different  elements,  it  is  evident  that  the  ratio 
between  the  weights  of  the  elements  in  a  compound  must  be 
the  same  as  the  ratio  between  the  weights  of  the  atoms  composing 
the  smallest  particle  of  the  compound.  Thus  if  the  smallest 
particle  (molecule)  of  water  which  can  exist  contains  two  atoms 


INTRODUCTION  15 

of  hydrogen  united  to  one  atom  of  oxygen  and  an  atom  of  oxygen 
weighs  16  times  as  much  as  an  atom  of  hydrogen,  any  quantity 
of  water,  whether  large  or  small,  must  contain  hydrogen  and 
oxygen  in  the  proportion  of  two  to  sixteen.  If,  for  instance, 
1000  atoms  of  oxygen  could  be  mixed  with  2001  atoms  of  hydro- 
gen, after  combination  had  taken  place  one  atom  of  hydrogen 
would  be  left  uncombined.  In  this  way  the  theory  explains 
very  satisfactorily  the  law  of  constant  proportion.  It  explains 
equally  well  the  law  of  combining  weights,  for  these  combining 
weights  must  be  directly  connected  with  the  relative  weights 
of  the  atoms  of  the  elements. 

In  accordance  with  the  atomic  theory  we  may  select  some  ele- 
ment as  our  unit  for  atomic  weights,  and  by  determining  the 
amounts  of  other  elements  which  combine  with  a  given  weight  of 
this  element  and  the  number  of  atoms  of  each  element  in  the 
compounds  formed,  we  can  determine  the  weights  of  the  atoms 
of  the  other  elements  as  compared  with  the  weight  of  an  atom 
of  the  element  taken  as  a  unit.  Thus  if  we  take  hydrogen  as  our 
unit  and  find  that  hydrochloric  acid  contains  one  part  of  hydro- 
gen to  35.5  parts  of  chlorine,  and  can  show,  further,  that  a  mole- 
cule of  hydrochloric  acid  contains  one  atom  of  chlorine  and  one 
atom  of  hydrogen  (p.  92),  the  atom  of  chlorine  must  be  35.5 
times  as  heavy  as  the  atom  of  hydrogen  and  we  say  that  the 
atomic  weight  of  chlorine  is  35.5.  Or  if  we  find  that  water  con- 
tains 8  parts  of  oxygen  for  one  of  hydrogen  1  and  a  molecule  of 
water  contains  one  atom  of  oxygen  and  two  atoms  of  hydrogen, 
the  atom  of  oxygen  must  be  16  times  as  heavy  as  the  atom  of 
hydrogen  and  we  say  that  the  atomic  weight  of  oxygen  is  16. 

The  atomic  theory,  which  could  be  considered*  as  scarcely 
more  than  a  doubtful  hypothesis  when  it  wag  first  proposed  by 
Dalton,  became  the  central,  guiding  principle  in  the  development 
of  the  science  of  chemistry  during  the  nineteenth  century ;  and 
evidence  in  its  favor  has  been  accumulated  from  very  many  dif- 
ferent and  independent  directions,  so  that,  now,  the  actual 
existence  of  atoms  and  molecules  can  scarcely  be  doubted. 
1  The  exact  composition  of  water  will  be  considered  later. 


16  A  TEXTBOOK  OF  CHEMISTRY 

We  even  have  a  half  dozen  different  ways  of  estimating  the 
actual  weight  of  an  atom  and  the  estimates  agree  fairly  well. 
These  estimates  give  the  number  of  molecules  in  a  cubic  centimeter 
of  air  under  standard  conditions  as  about  2.71  X  1019  or  nearly 
thirty  million  million  millions  (Millikan.  See  also  Rutherford, 
Presidential  Address  before  Section  A  of  the  British  Association 
at  the  Winnipeg  meeting).  Sir  William  Thomson  (known  later 
as  Lord  Kelvin)  once  used  the  illustration  that  if  a  drop  of  water 
could  be  magnified  to  the  size  of  the  earth  the  molecules  would 
be  larger  than  small  shot  and  smaller  than  cricket  balls.  This 
is  something  the  same  sort  of  an  estimate  as  if  we  were  to  say  that 
a  certain  animal  is  the  size  of  a  dog.  Our  knowledge  of  the  space 
filled  by  a  molecule  is  now  much  more  accurate. 

Selection  of  Atomic  Weights.  In  the  series  of  compounds 
used  to  illustrate  the  law  of  combining  weights,  the  combining 
weights  of  oxygen  are  8,  16  and  48.  If  the  table  were  extended, 
the  values  4  and  32  might  be  found  in  other  compounds,  and 
almost  any  multiple  of  8.  It  is  evident  that  if  we  start  with 
hydrogen  and  give  it  an  atomic  weight  of  1  (see,  however,  p.  72), 
only  one  of  these  various  combining  weights  can  be  the  true 
atomic  weight  of  oxygen.  Since  the  atoms  and  molecules  are 
so  small  as  to  be  beyond  the  possibility  of  direct  observation,  it 
seemed  for  a  long  time  impossible  to  select  the  true  atomic  weight 
from  among  the  various  possible  combining  weights.  Dal  ton 
thought  it  most  natural  to  suppose  that  the  molecule  of  water 
contains  one  atom  of  hydrogen  and  one  atom  of  oxygen  and  on 
this  basis  the  atomic  weight  of  oxygen  would  be  8  instead  of  16. 
The  reasons  for  considering  that  the  true  atomic  weight  of  oxygen 
is  16  and  thte  methods  used  in  selecting  what  are  believed  to  be 
true  atomic  weights  will  be  considered  later  (p.  92). 

Formulas.  The  atomic  weights  selected  for  the  elements  used 
to  illustrate  the  law  of  combining  weights  are:1  H  =  1, 
Cu  =  63.6,  S  =  32,  Cl  =  35.5;  Fe  =  56;  Na  =  23.  If  we 

1  These  values  are  rounded  off.  The  accepted  values  are : 
H  =  1.008,  O  =  16.00,  Cu  =  63.57,  S  =  32.07,  Cl  =  35.46,  Fe  = 
55.84,  Na  =  23.00. 


INTRODUCTION  17 

express  the  composition  of  water,  cuprous  oxide  and  cupric 
sulfide  in  such  a  manner  as  to  avoid  the  use  of  fractions  of  atomic 
weights,  the  ratios  for  these  compounds  become : 

Water,  H :  O   =       2:16 

Cuprous  oxide,    O  :  Cu  =      16  :  127.2 
Cupric  sulfide,  Cu  :  S     =  63.6  :  32 

In  accordance  with  the  atomic  theory  it  follows  from  these 
ratios  that  a  molecule  of  water  contains  two  atoms  of  hydrogen 
for  each  atom  of  oxygen,  that  a  molecule  of  cuprous  oxide  con- 
tains two  atoms  of  copper  for  one  of  oxygen  and  cupric  sulfide 
contains  the  same  number  of  atoms  of  sulfur  as  of  copper  in  its 
molecule.  It  has  been  found  very  convenient  to  express  these 
relations  by  using  the  symbol  of  each  element  to  stand  for  one 
atom  of  the  element  and  so  to  write  formulas  for  compounds, 
using  numerical  subscripts  to  designate  the  number  of  atoms  of 
each  element  contained  in  a  molecule  of  the  compound.  The 
formulas  for  the  compounds  are:  H2O,  Cu2O,  CuS.  Since  a 
formula  is  always  based  on  the  proportion  by  weight  of  each 
element  contained  in  the  compound,  it  tells  us  not  only  how 
many  atoms  of  each  element  are  contained  in  a  molecule  of  the 
compound,  but  it  also  tells  us  the  exact  composition  of  the  com- 
pound by  weight.  Thus  the  formula  H2SO4,  for  sulfuric  acid, 
means  that  a  molecule  of  sulfuric  acid  contains  two  atoms  of 
hydrogen,  one  atom  of  sulfur,  and  four  atoms  of  oxygen ;  but 
it  also  means  that  the  acid  is  composed  of  2  parts  by  weight 
of  hydrogen,  32  parts  of  sulfur  and  64  parts  of  oxygen. 

Strictly  speaking,  the  formulas  given  should  be  written  H2Oi, 
Cu2Ob  CuiSi  and  H2SiO4,  but  by  common  consent  the  subscript 
1  is  always  understood  when  no  subscript  is  given. 

What  are  the  formulas  of  the  other  compounds  mentioned  on 
p.  14? 

Composition  of  Pure  Substances.  From  what  has  been  stated 
we  may  derive  a  fifth  characteristic  of  a  pure  substance.  The 
composition  of  a  pure  substance  can  always  be  expressed  by 
exact  multiples  of  the  atomic  weights  of  the  elements  compos- 


18  A  TEXTBOOK  OF  CHEMISTRY 

ing  it.  This  may  be  considered  as  still  another  way  of  stating 
the  law  of  combining  weights.  The  law  has  been  tested  by  the 
analysis  of  thousands  of  compounds,  and,  like  the  law  of  constant 
proportion,  it  is  one  of  the  perfect  laws  from  which  no  deviation 
has  been  discovered. 

Study  of  Chemistry.  To  obtain  a  knowledge  of  the  elements 
of  chemistry  it  is  necessary  to  become  acquainted  with  a  large 
number  of  facts  about  the  substances  with  which  the  science 
deals,  but  it  is  still  more  important  to  understand  the  relations 
connecting  these  facts  with  each  other  and  the  fundamental 
laws  and  theories  by  which  the  facts  are  grouped  together  and 
explained.  Success  in  the  study  depends  especially  on  the  abil- 
ity to  learn  new  facts  in  their  relation  to  those  which  have  al- 
ready been  acquired  and  on  the  cultivation  of  a  logical  as  dis- 
tinguished from  an  arbitrary  memory.  Formulas,  especially, 
should  be  derived,  whenever  possible,  from  the  formulas  of  other 
compounds  of  the  same  elements,  and  not  learned  individually, 
except  in  the  earliest  portion  of  the  study.  Chemical  equations 
should  be  written  on  the  basis  of  a  knowledge  of  the  reacting 
substances  and  of  the  products  of  the  reaction,  and  should  never 
be  learned  by  brute  memory. 

In  the  systematic  treatment  of  the  subject  the  more  common 
elements  will  be  considered  first  and  under  each  element  the 
compounds  of  that  element  with  each  of  those  previously  studied 
will  be  mentioned  so  far  as  this  is  desirable.  See  p.  132. 


CHAPTER  II 
OXYGEN 

SYMBOL,  O.    ATOMIC  WEIGHT,  16. 

Occurrence.  Oxygen  is  the  most  abundant  and  one  of  the 
most  important  of  all  the  elements.  It  forms  about  one  fifth 
of  the  volume  of  the  air,  eight  ninths  of  the  weight  of  water  and 
nearly  one  half  the  weight  of  the  mineral  substances  which  com- 
pose the  crust  of  the  earth.  Oxygen  is  found  in  all  living  bodies 
and  is  a  constituent  of  a  larger  number  of  compounds  than  any 
other  element  except  carbon. 

Preparation.  1.  When  metallic  mercury  is  heated  at  the 
right  temperature  in  contact  with  the  air,  it  is  slowly  converted 

B 


G 


Fig.  2 

into  a  bright  red  compound  called  oxide  of  mercury.  The  French 
chemist,  Lavoisier,  carried  out  the  experiment  in  the  apparatus 
shown  in  the  figure,  and  proved  that  after  the  mercury  had  been 
heated  several  weeks  the  air  no  longer  decreased  in  volume  and 
he  concluded  that  this  was  because  the  oxygen  of  the  air  had 

19 


20  A  TEXTBOOK  OF  CHEMISTRY 

all  been  removed  by  combination  with  the  mercury.  He  then 
collected  the  oxide  of  mercury  and  heated  it  to  a  higher  tempera- 
ture till  it  was  all  decomposed  into  mercury  and  gaseous  oxygen. 
The  volume  of  oxygen  was  the  same  as  the  decrease  in  volume 
of  the  air  during  the  heating  in  contact  with  the  mercury. 

The  quantitative  relation  between  the  mercury,  oxygen  and 
oxide  of  mercury  may  be  very  briefly  expressed  by  means  of  the 
symbols  for  the  elements,  as  follows : 

Hg  +  O  HgO 

Mercury  Oxygen  Oxide  of  Mercury 

HgO  Hg         +  O 

Oxide  of  Mercury  Mercury  Oxygen 

Since  the  atomic  weight  of  mercury  is  200  and  the  atomic 
weight  of  oxygen  is  16,  the  first  equation  means  that  200  parts 
by  weight  of  mercury  combine  with  16  parts  by  weight  of  oxy- 
gen to  form  216  parts  of  oxide  of  mercury ;  and  the  second  equa- 
tion means  that  216  parts  of  oxide  of  mercury  decompose  into 
200  parts  of  mercury  and  16  parts  of  oxygen.  It  will  be  noticed 
that  the  symbols  of  two  elements  placed  side  by  side  represent 
a  compound,  while  a  symbol  by  itself  represents  a  free 
element. 

2.  The  portion  of  air  which  was  not  absorbed  by  the  mercury 
was  chiefly  nitrogen  and  formed  about  four  fifths  of  its  volume. 
Liquid  nitrogen  boils  at   —  194°,  while  liquid  oxygen  boils  at 
—  182.5°.     If  liquid  air  is  allowed  to  boil,  the  nitrogen  goes  off, 
chiefly,  at  first,  and  the  gas  which  comes  off  toward  the  end  is 
nearly  pure  oxygen.     In  this  way  the  oxygen  and  nitrogen  may 
be  separated  very  much  as  alcohol  and  water  are  separated  by 
distillation.     Oxygen  prepared  in  this  manner  is  compressed  into 
strong  steel  cylinders  for  medicinal  and  other  uses. 

3.  When  potassium  chlorate  is  heated,  it  melts  and  begins 
to  decompose    slowly    into    potassium    chloride    and    oxygen. 
Potassium    chlorate   has    the  composition  represented  by  the 
formula  KC1O3. 


OXYGEN  21 

Potassium,  K  =    39.1    parts  or    31.90  per  cent 

Chlorine,     Cl  =    35.46  parts  or    28.93  per  cent 

Oxygen,    3  O  =    48.      parts  or    39.17  per  cent 

Total  122.56  parts  or  100.      per  cent 

The  decomposition  may  be  represented  by  the  equation : 
KC1O3    =        KC1        +      3O 

Potassium          Potassium  Oxygen 

Chlorate  Chloride 

If  the  potassium  chlorate  is  mixed  with  one  fourth  of  its 
weight  of  finely  powdered  manganese  dioxide,  MnO2,  the  de- 
composition will  begin  at  a  much  lower  temperature  and  pro- 
ceed more  rapidly  than  when  the  potassium  chlorate  is  heated 
alone.  If  the  residue  in  the  retort,  after  the  decomposition  is 
complete,  is  treated  with  water,  the  potassium  chloride  will 
dissolve,  while  the  manganese  dioxide  will  remain  undissolved 
and  may  be  readily  separated  from  the  solution  by  filtration. 
If  the  manganese  dioxide  is  examined,  it  will  be  found  that  it 
has  not  changed  in  composition  or  amount.  We  may,  there- 
fore, write  the  equation : 

KC103     +      Mn02      =       KC1      +    3O    +      MnO2 

Potassium  Manganese  Potassium        Oxygen         Manganese 

Chlorate  Dioxide  Chloride  Dioxide 

4.  When  fused  sodium  peroxide  containing  a  very  little 
copper  oxide1  is  dissolved  in  water,  it  gives  sodium  hydroxide 
and  oxygen : 

Na202      +      H2O      =      2NaOH      +       O 

Sodium  Water  Sodium  Oxygen 

*    Peroxide  Hydroxide2 

As  the  copper  oxide  is  left  unchanged  at  the  end  and  as  the 
reaction  will  take  place,  though  more  slowly,  in  its  absence,  we 
may  omit  it  in  writing  the  equation. 

1  The  substance  is  known  commercially  as  "  oxone." 

2  The  student  should  notice  the  connection  between  the  name  and 
the  composition  of  sodium  hydroxide.     Many  similar  compounds 
containing  oxygen  and  hydrogen  are  called  hydroxides. 


22 


A  TEXTBOOK  OF  CHEMISTRY 


Collection  and  Storage  of  Gases.  Gases  which  are  only 
slightly  soluble  in  water  and  which  are  not  required  in  a  high 
state  of  purity  are  conveniently  collected  and 
stored  in  a  gasometer  of  the  form  shown  in 
Fig.  3.  If  the  stopcocks  A  and  B  are  opened 
and  the  cap  C  screwed  on,  water  placed  in  the 
cup  above  will  run  into  the  body  of  the  gas- 
ometer till  it  is  filled  with  water  and  all  of 
the  air  is  expelled.  Now  on  closing  the  stop- 
cocks the  gasometer  will  still  remain  filled  with 
water  after  removing  the  cap  at  C.  By  means 
of  a  tube  inserted  through  C  gas  may  be  in- 
troduced and  will  fill  the  gasometer,  displacing 
the  water,  which  will  flow  out  of  C  by  the 
side  of  the  tube  delivering  the  gas.  After 
filling  the  gasometer  and  replacing  the  cap  at 
C,  on  opening  the  stopcock  A,  water  will  run 
into  the  body  of  the  gasometer  and  the  gas 
may  be  drawn  off  as  desired  through  B. 

Properties  of  Oxygen.  Oxygen  is  a  colorless, 
odorless  and  tasteless  gas.1  The  weight  of  one  liter  at  0° 
and  under  a  pressure  of  760  millimeters  of  mercury  (about 
the  average  atmospheric  pressure  at  sea  level)  is  1.429  grams.2 
Under  the  same  conditions  of  temperature  and  pressure  it  is 
about  one  tenth  heavier  than  the  same  volume  of  air. 

The  most  striking  property  of  oxygen  is  the  vigor  with 
which  it  supports  combustion.  All  substances  which  burn  in 
air  burn  much  more  rapidly  and  vigorously  in  oxygen.  A  splin- 
ter of  wood  having  a  live  coal  on  the  end  will  burst  into  flame, 

1  These  statements  refer,  of  course,  to  the  ordinary  form  of  the 
element.     Ozone,  which  is  another  form  of  oxygen,  is  colored  and  has 
a  strong  odor  (p.  98). 

2  At  45°  latitude.    It  is  slightly  less  at  lower  latitudes  because  the 
pressure  of  760  mm.  of  mercury  is  less  and  the  density  of  the  gas  is 
less.    At  the  latitude  of  New  York  the  weight  of  one  liter  of  oxygen 
is  1.42845  grams.     If,  however,  the  reading  of  the  barometer  is  cor- 
rected for  latitude  and  altitude,  the  weight  of  one  liter  of  oxygen  is 
1.429  grams  at  any  place. 


Fig.  3 


OXYGEN 


23 


if  thrust  into  the  gas.  A  piece  of  charcoal,  barely  ignited,  will 
glow  intensely  and  be  surrounded  by  a  pale  blue  flame,  scarcely 
visible  in  the  intense  light  of  the  glowing  mass. 
The  product  of  the  combustion  is  carbon  di- 
oxide, a  colorless  gas.  Sulfur  burns  with  a 
brilliant  blue  flame,  giving  sulfur  dioxide,  also 
a  colorless  gas.  Phosphorus  burns  with  an 
intense  white  light,  giving  a  white,  solid  com- 
pound, phosphorus  pentoxide.  A  coil  of  iron 
wire  or  a  steel  watch  spring  to  which  is  at- 
tached a  string  that  has  been  dipped  in  melted 
paraffin,  may  be  set  on  fire  and  will  burn  in 
oxygen  (Fig.  4),  throwing  off  brilliant  sparks 
and  forming  white-hot,  molten  globules  of  the 
magnetic  oxide  of  iron,  which  will  drop  off 
from  time  to  time. 


Fig.  4 


The  equations  which  represent  the  quantitative  relations  in 
these  experiments  are : 

=  CO2 

Carbon  Dioxide 

=  SO2 

Sulfur  Dioxide 

=  P205 

Phosphorus  Pentoxide 

=  Fe304 

Magnetic  Oxide  of  Iron 

Oxygen  and  Acid  Properties.  If  sulfur  dioxide  or  phosphorus 
pentoxide  is  dissolved  in  water,  the  solution  obtained  will  have 
a  sour  taste,  and  acid  properties.  Many  other  compounds  of 
nonmetallic  elements  with  oxygen  combine  with  water  in  a 
similar  manner  to  form  acids,  and  it  is  because  of  this  that  the 
name  oxygen,  meaning  "  acid  former,"  was  first  given  to  the 
element.  When  the  name  was  given,  it  was  supposed  that  all 
acids  contain  oxygen,  but  it  was  discovered  later  that  this  is 
not  the  case. 


c 

4-        2O 

Carbon 

Oxygen 

s 

+       20 

Sulfur 

Oxygen 

2P 

+      5O 

Phosphorus 

Oxygen 

3Fe 

+      40 

Iron 

Oxygen 

24 


A  TEXTBOOK  OF  CHEMISTRY 


Combustion.  Effect  of  Concentration  on  a  Chemical  Reac- 
tion. The  similarity  between  ordinary  combustion  and  the 
burning  of  substances  in  oxygen  is  apparent.  A  more  careful 
study  of  the  matter  has  shown  that  exactly  the  same  compounds 
are  formed  when  charcoal,  sulfur  or  phosphorus  burn  in  the  air 
as  are  formed  when  they  burn  in  oxygen,  and  even  in  the  case  of 
iron,  the  magnetic  oxide  formed  by  burning  the  steel  watch 
spring  has  just  the  same  composition  as  that  of  the  scale  formed 
when  white-hot  iron  is  exposed  to  the  air.  The  burning  of  iron 

in  air  can  also  be  shown  by  sprinkling 
fine  iron  filings  through  a  flame. 
About  four  fifths  of  the  air  by  volume 
is  nitrogen.  This  does  not  combine  with  the 
burning  substances,  and  by  its  presence  it 
moderates  the  action  of  the  oxygen,  partly 
by  diluting  it,  partly  because  it  must  be 
heated  to  the  same  temperature  as  the  other 
substances,  and  this  absorbs  a  large  part  of 
the  heat  of  the  reaction  and  so  lowers  the 
temperature  to  which  the  burning  substance 
is  heated.  These  facts  illustrate  two  prin- 
ciples of  almost  universal  application  in 
chemistry :  first,  that  the  speed  of  a  chemi- 
cal reaction  is  increased  by  increasing  the 
concentration  of  one  of  the  reacting  sub- 
stances, here  the  concentration  of  the  oxygen;  and,  second, 
that  the  speed  of  a  reaction  is  affected  by  the  temperature  and 
is  greater  at  high  temperatures  than  at  low  ones.  We  shall  find 
later  that  the  first  of  these  principles  can  be  stated  in  the  form 
of  an  accurate,  quantitative  law,  but  the  phenomena  of  ordi- 
nary combustion  are  not  well  suited  for  a  quantitative  study 
of  this  kind. 

Kindling  Temperature.  If  we  place  on  an  iron  plate  smail 
pieces  of  phosphorus,  sulfur  and  charcoal,  it  will  be  found  on 
warming  the  plate  that  phosphorus  takes  fire  at  a  quite  low 
temperature,  the  sulfur  at  a  moderate  heat,  while  the  charcoal 


Fig.  5 


CALORIMETER 


25 


will  not  burn  till  the  plate  is  nearly  red-hot.  The  temperature 
at  which  combination  of  a  substance  with  the  oxygen  of  the  air 
is  sufficiently  rapid  so  that  it  takes  fire  is  called  the  kindling 
temperature.  The  temperature  rises  rapidly  from  the  heat  of  the 
reaction  as  soon  as  the  substance  is  kindled.  Well-known  appli- 
cations of  the  gradations  of  kindling  temperature  are  the  old- 
fashioned  sulfur  match  and  the  methods  commonly  used  in  kin- 
dling a  fire.  Kindling  temperature  is  not  a  satisfactory  measure 
of  the  affinity  of  a  substance  for  oxygen  nor  is  it  closely  con- 
nected with  the  heat  generated  on  combination  with  oxygen. 

Heat  of  Combustion.  Calorimeter.  When  substances  burn, 
a  part  of  the  chemical  energy  of  the  burning  substance  and  of 
the  chemical  energy  of  the  oxygen  is  converted  into  heat.  The 
amount  of  energy  transformed  into 
heat  when  one  gram  of  the  substance 
burns,  or,  for  most  scientific  pur- 
poses, the  energy  obtained  from  one 
gram  atom  or  gram  molecule  of  the 
substance,  is  called  its  heat  of  com- 
bustion. By  gram  atom  is  meant  as 
many  grams  of  the  substance  as 
there  are  units  in  the  atomic  weight 
—  as  12  grams  of  carbon  or  31  grams 
of  phosphorus.  By  gram  molecule  is 
meant  as  many  grams  of  the  sub- 
stance as  there  are  units  in  the  mo- 
lecular weight,  as  44  grams  of  carbon 
dioxide  (12  grams  of  carbon  -f-  32 
grams  of  oxygen). 

The  heat  of  combustion  is  de- 
termined in  an  instrument  called  a 
calorimeter,  the  most  accurate  form 
being  known  as  a  bomb- calorimeter 
because  the  combustion  is  carried 


Fig. 


out  in  an  atmosphere  of  compressed  oxygen  in  a  strong,  tightly 
closed  bomb,  which  is  immersed  in  water.     The  substance  is 


26  A  TEXTBOOK  OF  CHEMISTRY 

placed  in  a  small  crucible  within  the  bomb  and  is  ignited  by 
means  of  a  fine  iron  wire,  which  is  heated  for  a  moment  by  an 
electric  current.  The  weight  of  the  substance,  the  weight  of 
water  surrounding  the  bomb  and  the  temperature  of  the  water 
before  and  after  the  substance  is  burned  are  accurately  de- 
termined. There  are,  of  course,  many  other  details  about  the 
apparatus  and  manipulation  which  need  not  be  described  here. 
(See  Atwater,  J.  Am.  Chem.  Soc.  25,  659,  and  Richards  and 
Jesse,  ibid.  32,  268.) 

The  amount  of  heat  required  to  raise  the  temperature  of  one 
gram  of  water  one  degree  at  15°  is  called  a  calorie  (see  p.  33). 
This  is  often  called  the  small  calorie  and  designated  by  the 
abbreviation  cal.  to  distinguish  it  from  the  large  Calorie,  which 
is  the  amount  of  heat  required  to  raise  the  temperature  of  a 
kilogram  of  water  one  degree  and  which  is  designated  by  the  ab- 
breviation Cal. 

*  The  corresponding  unit  of  the  English  system  is  the  British 
Thermal  Unit  (B.  T.  U.),  the  heat  required  to  raise  the  tempera- 
ture of  a  pound  of  water  1°  F.  Since  1  Kg.  =  2.204  Ib.  and  1°  F. 
=  |  of  1°  C.,  1  Cal.  =  3.968  B.  T.  U.  Heat  of  combustion  ex- 
pressed in  calories  per  kilogram,  however,  is  reduced  to  British 
Thermal  Units  per  pound  by  multiplying  by  -§• .  From  the  re- 
sults of  determinations  with  the  calorimeter  it  is  possible  to  cal- 
culate how  many  grams  of  water  can  be  raised  one  degree  in  tem- 
perature by  burning  one  gram  or  one  gram  atom  of  a  substance 
and  this  will  be  the  heat  of  combustion  of  the  substance  in 
calories.1 

The  heats  of  combustion  for  the  substances  mentioned  in  this 
chapter  are : 


1  Since  the  amount  of  heat  required  to  raise  the  temperature 
of  one  gram  of  water  one  degree  varies  slightly  with  the  tempera- 
ture, it  is  necessary  in  accurate  work  to  define  the  temperature 
at  which  the  experiment  is  performed.  A  mean  temperature  of 
15°  is  most  often  used.  On  account  of  the  variability  of  the 
calorie  it  has  been  proposed  to  use  the  joule  as  a  unit.  The  small 
calorie  is  equal  to  4.182  joules  at  15°.  See  G.  N.  Lewis,  Journal  of 
the  American  Chemical  Society,  35,  4  (1913). 


CHEMICAL  ENERGY  27 

For  one  gram  of  carbon  (charcoal),  8080  calories 
For  one  gram  of  sulfur  (rhombic),  2190  calories 
For  one  gram  of  phosphorus  (yellow),  5970  calories 
For  one  gram  of  iron  (to  Fe3O4),  1616  calories 

For  one  gram  of  mercury  (to  HgO),        105  calories 

For  one  gram  atom  of  carbon  (charcoal),         96,960  calories  1 
For  one  gram  atom  of  sulfur  (rhombic),  70,180  calories 

For  one  gram  atom  of  phosphorus  (yellow),  185,000  calories 
For  one  gram  atom  of  iron  (to  Fe3O4),  90,200  calories 

For  one  gram  atom  of  mercury  (to  HgO),       21,000  calories 

The  Nature  of  Chemical  Energy.  The  amount  of  energy 
liberated  by  a  burning  substance  is  very  large.  If  it  were 
possible  completely  to  transform  the  energy  liberated  by  burning 
a  pound  of  good  coal  into  mechanical  energy,  it  would  lift  a  ton 
weight  over  4000  feet.  It  is  a  very  good  engine  which  will  con- 
vert ten  per  cent  of  the  energy  of  the  coal  burned  under  its  boiler 
into  useful  work,  but,  in  spite  of  the  fact  that  more  than  ninety 
per  cent  of  the  energy  of  the  coal  is  dissipated  and  lost,  the  total 
energy  is  so  enormous  that  the  steam  engine  is  commercially 
economical. 

A  very  natural  question  which  arises  here  is;  What  is  the 
source  of  the  energy  which  suddenly  appears  as  heat  when  par- 
ticles of  carbon  and  oxygen  unite  to  form  carbon  dioxide  ?  Is 
there  some  motion  within  the  particles  of  oxygen  and  carbon 
which  is  transformed  into  heat  when  they  unite?  or  do  the 
particles  collide  when  their  mutual  attraction  brings  them  to- 
gether, somewhat  as  a  meteor  collides  with  the  earth?  For 
these  questions  there  are,  at  present,  no  answers,  and  specula- 
tions about  them  are  of  very  little  value  till  some  one  can  dis- 
cover some  sort  of  experimental  evidence  bearing  upon  them. 

1  In  joules  these  are  : 

For  one  gram  atom  of  carbon  (charcoal)  to  CO2,  405,700  joules 
For  one  gram  atom  of  phosphorus,  to  PzO&,  773,000  joules 

For  one  gram  atom  of  sulfur,  to  SO2,  293,000  joules 

For  one  gram  atom  of  iron,  to  Fe3O4,  3 77,000  joules 

For  one  gram  atom  of  mercury,  to  HgO,  87,800  joules 


28  A  TEXTBOOK  OF  CHEMISTRY 

It  is  well,  however,  to  recognize  how  imperfect  and  fragmentary 
our  knowledge  is  and  that  there  are  hundreds  of  questions  like 
these  for  which  we  have  no  answer.  It  is  also  well,  at  times, 
to  ask  such  questions  and  consider  whether  there  is  any  tangible 
method  of  attacking  the  problem,  for,  while  the  explanation  seems 
beyond  our  grasp,  at  present,  many  similar  problems  which 
would  have  seemed  beyond  the  possibility  of  a  solution  one  hun- 
dred years  ago  have  been  solved. 

Catalysis.  It  has  been  pointed  out  that  when  manganese 
dioxide  is  mixed  with  potassium  chlorate  the  latter  decomposes 
at  a  lower  temperature  or  more  rapidly  than  when  the  chlorate 
is  heated  by  itself,  but  that  the  manganese  dioxide  is  left  un- 
changed in  the  end.  When  a  substance  acts  in  this  manner  by 
its  mere  presence,  causing  a  reaction  or  decomposition  to  take 
place  at  a  lower  temperature  or  more  rapidly,  it  is  called  a 
catalytic  agent  and  the  process  is  called  catalysis. 

*  A  study  of  this  particular  case  makes  it  seem  probable  that 
the  oxygen  is  at  first  transferred  from  the  potassium  chlorate  to 
the  manganese  dioxide  and  that  the  compound  of  the  manganese 
dioxide  with  the  oxygen  decomposes  at  a  lower  temperature 
than  the  potassium  chlorate.  Such  an  explanation  seems,  at 
first,  paradoxical,  for  it  seems  to  imply  that  manganese  dioxide 
has  a  greater  affinity  for  oxygen  than  potassium  and  chlorine 
have  and  so  can  take  the  oxygen  away  from  the  potassium 
chlorate,  while  in  the  resulting  compound  the  affinity  of  the 
manganese  dioxide  for  oxygen  seems  to  be  less  than  that  of 
potassium  and  chlorine  for  oxygen,  because  the  decomposition 
of  the  manganese  compound  occurs  at  a  lower  temperature 
than  that  required  for  the  decomposition  of  potassium  chlorate. 

A  partial  explanation  of  this  seeming  paradox  is  found  in  the 
fact  that  very  many  substances  act  upon  each  other  chemically 
at  a  much  lower  temperature  than  that  at  which  either  decom- 
poses into  its  elements.  It  is  also  true  that  the  stability  of  a 
compound  is  not  an  accurate  measure  of  the  affinity  between  the 
elements  of  which  it  is  composed.  The  affinity  between  the 
elements  of  a  compound  which  decomposes  at  200°  is  not  neces- 


CHEMICAL  AFFINITY.    NOMENCLATURE  29 

sarily  less  than  that  between  the  elements  of  a  compound  which 
decomposes  at  400°.  Still  further,  it  is  by  no  means  always  true 
that  when  an  element  is  transferred  from  one  compound  to 
another  its  affinity  for  the  element  with  which  it  combines  is 
greater  than  that  for  the  element  which  it  leaves.  These  ques- 
tions will  be  considered  further  later ;  but  it  is  well,  at  the  outset, 
to  avoid  certain  misconceptions  which  are  very  liable  to  arise 
.  because  the  facts  of  chemistry  are  often  so  very  different  from 
what  our  first  and  most  natural  idea  of  chemical  affinity  would 
lead  us  to  expect. 

Chemical  Affinity.  The  term  affinity  has  been  used  in  the  pre- 
ceding paragraph  and  seems  to  call  for  some  definition.  The 
word  is  generally  used  in  a  rather  indefinite  way  to  designate 
that  attraction  between  elements  which  causes  them  to  unite  to 
form  compounds. 

*  The  real  nature  of  chemical  affinity  is  not  known.  This  is 
another  of  those  questions,  like  the  cause  of  the  heat  generated 
when  elements  combine,  which  waits  for  an  answer.  Doubtless 
the  two  questions  are  intimately  connected.  But  while  we  do 
not  know  its  real  nature,  we  can  learn  a  great  deal  about  the  con- 
ditions under  which  chemical  affinity  acts.  Thus  it  seems  al- 
ways, in  accordance  with  the  laws  of  constant  proportion  and  of 
combining  weights,  to  be  exerted  between  definite  quantities 
of  the  elements.  We  shall  find,  too,  that  there  are  certain  ways 
in  which  we  can  give  to  chemical  affinity  an  accurate,  mathemati- 
cal definition  and  measure  it  quantitatively. 

It  seems  natural  to  think  of  chemical  affinity  as  a  force  similar 
to  the  force  of  gravity  or  to  the  force  of  electrical  attraction. 
It  may  be  that  it  is  closely  connected  with  one  or  both  of  these. 

Nomenclature.  The  compounds  of  oxygen  have  been  called, 
in  this  chapter,  oxides.  This  is  an  application  of  a  system  of 
naming  substances  which  is  used  for  all  compounds  consisting 
of  two  elements.  As  compounds  of  oxygen  are  called  oxides, 
compounds  of  sulfur  are  called  sulfides,  compounds  of  chlorine, 
chlorides,  etc.  In  order  to  give  more  definite  names,  prefixes 
derived  from  the  Greek  numerals  are  used.  CO  is  called  carbon 


30  A  TEXTBOOK  OF  CHEMISTRY 

monoxide;  CO2,  carbon  dioxide;  HgCl2,  mercury  dichloride; 
P2O3,  phosphorus  trioxide ;  SO3,  sulfur  trioxide ;  CCU,  carbon 
tetrachloride ;  P2O5,  phosphorus  pentoxide.  In  addition  to 
these  names,  which  tell  how  many  atomic  weights  of  the  ele- 
ment are  contained  in  the  molecular  weight  of  the  substance, 
the  prefix  per  is  used  to  name  compounds  containing  more  oxy- 
gen than  some  other  oxide  of  the  same  element.  Thus  sodium 
peroxide,  Na2O2,  contains  more  oxygen  than  the  other  oxide  of 
sodium,  Na2O,  the  prefix  per  meaning  more  or  beyond. 

Still  another  method  of  naming  oxides  is  to  add  the  suffixes 
-ous  and  -ic  to  the  name  of  the  metal  or  other  element  which  is 
combined  with  oxygen.  The  ending  -ic  is  used  for  the  compound 
containing  the  larger  proportion  of  oxygen.  Thus  Hg2O  is  called 
mercurous  oxide;  HgO,  mercuric  oxide;  FeO,  ferrous  oxide; 
Fe2O3,  ferric  oxide. 

The  choice  among  these  three  methods  of  naming  oxides  and 
other  compounds  is  more  or  less  arbitrary  and  conventional. 


CHAPTER  III 
LAWS  OF  GASES 

UNITS  OF   LENGTH,  WEIGHT,  VOLUME,   TEMPERATURE,  TIME  AND 

ENERGY. 

Unit  of  Length.  Meter.  The  meter  was  originally  intended 
to  be  one  ten-millionth  of  the  distance  from  the  equator  to  the 
pole  of  the  earth,  measured  on  the  surface.  The  measurements 
by  means  of  which  the  first  meter  was  prepared  were  inaccurate, 
however,  and  the  real  meter  is  the  distance,  measured  at  the 
freezing  point  of  water,  between  two  marks  on  a  bar  of  platinum- 
iridium  kept  at  the  International  Bureau  of  Weights  and  Meas- 
ures at  Sevres,  France.  The  meter  is  divided  into  tenths,  hun- 
dredths  and  thousandths,  called  decimeters,  centimeters  and 
millimeters.  Its  most  common  multiple  is  the  kilometer, 
1000  meters. 

Unit  of  Weight.  Gram.  The  gram  was  intended  to  be  the 
weight  of  one  cubic  centimeter  of  water  at  its  maximum  density, 
4°  centigrade.  Here,  again,  the  first  measurements  were  not 
quite  accurate  and  the  real  kilogram  (1000  grams)  is  the  weight, 
in  a  vacuum,  of  a  mass  of  platinum-iridium  kept  at  the  Inter- 
national Bureau.  The  most  common  division  of  the  gram  is  one 
milligram,  the  thousandth  of  a  gram. 

Unit  of  Volume.  Liter.  The  liter  was  intended  to  have  a 
volume  of  one  cubic  decimeter.  Because  volumes  can  be  most 
accurately  compared  by  weighing  the  water  which  fills  them, 
the  real  liter  is  the  volume  occupied  by  one  kilogram  of  water, 
weighed  in  a  vacuum  at  4°  C.  The  cubic  centimeter  is  defined, 
conventionally,  as  one  one- thousandth  of  a  liter.1 

1  The  actual  weight  of  water  contained  in  a  cube  whose  edge  is 
one  centimeter  is  0.999982  g.  according  to  the  best  measurements. 
Because  the  edge  of  a  cube  of  water  weighing  one  gram  is  not  ex- 

31 


32  A  TEXTBOOK  OF  CHEMISTRY 

Units  of  Time.  The  units  of  time  used  in  chemistry  are  the 
second,  minute,  hour,  day  and  year.  These  are  all  fixed  by 
means  of  astronomical  observations  with  the  aid  of  accurate 
clocks. 

Unit  of  Temperature.  The  freezing  point  of  water  under  at- 
mospheric pressure  has  been  selected  as  zero  for  the  ordinary 
centigrade  scale  of  temperature,  and  the  boiling  point  of  water 
under  atmospheric  pressure  as  100°.  For  the  International  scale, 
the  interval  between  the  two  points  is  divided  into  one  hundred 
equal  parts  by  measuring  the  increase  in  pressure  of  hydrogen 
gas,  at  constant  volume,  the  initial  pressure  being  that  of  a 
column  of  mercury  one  meter  high  at  0°.  Absolute  tempera- 
tures will  be  considered  later. 

It  is  well  to  notice  that  the  unit  of  temperature  is  a  unit  of 
intensity  and  not  of  quantity.  In  this  respect  it  corresponds 
to  the  height  to  which  a  weight  is  raised  in  mechanical  energy 
or  to  the  volt  in  electricity. 

Units  of  Energy.  Kilogram-meter;  Erg.  The  simplest  unit 
of  energy  is  the  kilogram-meter,  the  energy  required  to  lift  one 
kilogram  to  a  height  of  one  meter.  Since  the  force  of  gravity 
varies  with  the  latitude  and  altitude,  another  unit,  which  is 
independent  of  these,  is  often  used.  This  is  the  erg,  and  is  twice 
the  energy  of  one  gram  1  moving  with  a  velocity  of  one  centimeter 
a  second.  Or  it  may  be  defined  as  the  energy  required  to  im- 
part to  one  gram  a  velocity  of  one  centimeter  per  second  or  to 
increase  its  velocity  by  one  centimeter  per  second.  One  joule  is 
10,000,000  ergs. 

actly  one  centimeter  in  length,  some  persons  prefer  to  call  the  con- 
ventional cubic  centimeter  a  mimliter.  The  suggestion  has  not, 
however,  been  generally  accepted. 

1  This  is  more  often  stated  as  the  mass  of  one  gram,  but  since 
weights  are  always  accurately  determined  by  the  balance,  one  gram 
determined  by  weighing  is  just  as  "absolute"  a  quantity  of  matter 
as  the  mass  of  one  gram.  For  the  same  reason  chemists  are  justi- 
fied in  speaking  of  atomic  weights  instead  of  atomic  masses.  It  is 
well  to  remember,  however,  that  weight  is  in  its  accurate,  scientific 
use,  a  measure  of  a  force  and  not  a  measure  of  a  quantity  of  matter. 
In  the  common  everyday  use  of  the  word  we  use  it  for  a  quantity 
of  matter. 


UNITS  33 

Centimeter-gram-second  System.  Absolute  units.  Since  in 
accordance  with  the  law  of  conservation  of  energy  every  form  of 
energy  bears  an  exact,  quantitative  relation  to  every  other,  any 
quantity  of  energy  which  we  can  measure  may  be  expressed  in 
terms  of  the  velocity  of  a  moving  mass.  The  units  necessary 
for  such  a  purpose  are  a  unit  of  length,  a  unit  of  mass  and  a 
unit  of  time.  Physicists  have  agreed  upon  the  centimeter, 
gram  and  second  as  fundamental  units  and  have  developed  a 
system  of  "  absolute  units  "  in  which  all  forms  of  energy  are 
measured  by  reference  to  these.  These  units  are  called  absolute 
because  they  are  independent  of  the  force  of  gravity. 

Units  of  Mechanical  Energy.  The  absolute  unit  for  mechan- 
ical energy  is  the  erg,  which  has  been  defined  above.  The  most 
common  unit  used  by  engineers  is  the  kilogram-meter  (or  the 
foot-pound  in  the  English  system).  At  45°  latitude  and  sea 
level  the  kilogram-meter  is  98,066,700  ergs. 

Unit  of  Power.  Power  is  the  rate  of  production  of  energy. 
One  horse  power  is  4600  kilogram-meters  or  33,000  foot-pounds 
per  minute. 

Units  of  Heat.  The  calorie  is  the  heat  required  to  raise  the 
temperature  of  a,  gram  of  water  one  degree  (p.  26).  It  varies 
slightly  with  the  temperature  and  for  accurate  work  the  tem- 
perature must  be  specified  —  usually  a  temperature  at  15° l  is 
taken  as  the  standard.  As  an  absolute  unit,  independent  of 
the  temperature  of  the  water  the  joule  has  been  suggested. 
One  calorie  at  15°  is  equal  to  4.187  joules. 

Electrical  Units.  The  primary  electrical  units  are  the  wit, 
ohm  and  ampere.  These  are  so  related  that  an  electromotive 

1  If  the  calorie  at  15°  is  taken  as  one,  the  values  of  the  calorie  at 
other  temperatures  are  as  follows  : 

10°  1.0016 

15°  1.0000 

20°  0.9991 

25°  0.9988 

30°  0.9989 

These  are  the  mean  of  the  values  of  Georg  lanke,  Ann.  Tables  of 
Physical  Constants  for  1910,  and  of  Bausfield,  Phil.  Trans.  211,  A, 
199  (1911). 


34  A  TEXTBOOK  OF  CHEMISTRY 

force  (E.  M.  F.)  of  one  volt  acting  through  a  resistance  of  one 
ohm  gives  a  current  of  one  ampere  or : 

Amperes  =    E'  M '  F ' 


R(in  ohms) 

The  unit  of  electrical  power  is  the  watt,  a  current  of  one  ampere 
flowing  under  a  difference  of  potential  of  one  volt.  It  is  equiva- 
lent to  10,000,000  ergs  or  one  joule  per  second.  The  kilowatt 
is,  of  course,  1000  watts  and  is  the  most  common  measure  for 
electrical  service  in  lighting,  heating,  running  of  motors  and  the 
like.  The  watt  is  one  of  the  "  absolute  "  units. 

An  electrical  horse  power  is  746  watts,  and  is  equivalent,  of 
course,  to  33,000  foot  pounds  per  minute. 

Chemical  Energy.  By  chemical  energy  is  meant  the  energy 
which  appears  as  mechanical  energy,  heat,  light,  sound  or  elec- 
tricity when  two  or  more  elements  unite,1  or  when  an  element  is 
changed  from  one  form  to  another,  as  ozone  to  oxygen.  It  is 
usually  expressed  in  terms  of  heat  units  or  electrical  units.  It 
must  always  refer  to  some  definite  chemical  action  which  takes 
place  and  can  never  refer  to  the  total  energy  contained  in  an  ele- 
ment or  compound,  as  we  have  no  means  of  measuring  this. 

Effect  of  Pressure  on  a  Gas.  Law  of  Boyle.  When  the  pres- 
sure applied  to  a  gas  is  doubled,  the  volume  is  reduced  to  one 
half ;  or  when  the  pressure  is  reduced  to  one  half,  the  volume  be- 
comes twice  as  great.  Another  method  of  stating  this  property 
of  gases  in  a  general  way  is  to  say  that  the  volume  of  a  gas  varies 
inversely  as  the  pressure.  Or,  mathematically  : 

V :  V  : :  Pf :  P,  or  VP  =  V'P'  =  Constant, 

where  V  and  V  are  two  volumes  of  the  same  quantity  of  gas  and 
P  and  Pf  are  the  corresponding  pressures.  This  is  known  as 
Boyle's  law.  It  is  not  an  accurate  law,  as  the  law  of  constant 

1  Heat  is  absorbed  when  some  substances  unite,  and  in  such  cases 
the  energy  of  the  compound  is  considered  as  negative  or  less  than 
nothing  in  comparison  with  that  of  the  elements  from  which  it  is 
formed,  but  the  idea  that  energy  can  be  really  negative  seems 
absurd. 


LAWS  OF  GASES  35 

proportion  is,  but  is  sufficiently  accurate  for  use  in  all  ordinary 
cases. 

*  The  extent  of  the  deviation  from  the  law  for  several  gases  is 
shown  in  the  following  table  : 

TABLE 

Volumes  filled  at  0°  by  two  liters  of  each  gas  when  the  pressure  is 
increased  from  one  atmosphere  to  two  atmospheres. 

Hydrogen  1.0006  liters 

Nitrogen  0.9996  liters 

Carbon  monoxide  0.9995  liters 

Oxygen  0.9991  liters 

Nitric  oxide  0.9989  liters 

Carbon  dioxide  0.9931  liters 

Nitrous  oxide  0.9924  liters 

Hydrochloric  acid  0.9919  liters 

Ammonia  0.9845  liters 

Sulfur  dioxide  0.9739  liters 

Those  gases  which  are  liquefied  most  easily  depart  farthest 
from  the  law,  and  all  gases  except  hydrogen  and  helium  are  com- 
pressed more  than  they  should  be  under  moderate  pressures. 
For  a  pressure  of  many  atmospheres  a  point  is  reached  where  all 
gases  which  do  not  liquefy  are  compressed  less  than  they  should 
be  in  accordance  with  the  law.  According  to  the  kinetic  theory 
(p.  58)  the  greater  compressibility  under  moderate  pressure 
is  caused  by  the  attraction  of  the  molecules  for  each  other  — 
the  same  forces  which  cause  the  gas  to  liquefy  at  low  tempera- 
tures or  under  pressure.  The  point  of  too  little  compressibility  is 
reached  when  the  molecules  are  brought  so  close  together  that  the 
molecules  themselves  fill  a  considerable  fraction  of  the  total  space. 

The  law  may  be  easily  illustrated  by  taking  a  gas  measuring 
tube,  graduated  in  cubic  centimeters,  filling  it  partly  full  of 
mercury  and  immersing  the  mouth  in  a  deep,  narrow  jar  con- 
taining mercury.  It  is  evident  that  if  the  tube  is  raised  or 
lowered  till  the  top  of  the  mercury  within  the  tube  is  exactly 
level  with  the  surface  of  the  mercury  on  the  outside,  the  pressure 


36 


A  TEXTBOOK  OF  CHEMISTRY 


of  the  gas  within  the  tube  will  be  the  same  as  that  shown  by 
a  barometer  in  the  same  room.1  If,  now,  the  tube  is 
raised,  the  volume  of  the  gas  will  be  seen  to  in- 
crease, and  for  any  given  position  the  pressure  of 
the  gas  must  be  equal  to  the  reading  of  the  barome- 
ter less  the  height  of  the  mercury  in  the  tube  above 
that  in  the  jar.  By  reading  the  volumes  in  two 
different  positions  of  the  tube  and  determining  the 
corresponding  pressures  the  data  for  a  verification  of 
the  law  may  be  easily  obtained. 

For  practical  uses  it  is  convenient  to  select  some 
standard  pressure  to  which  the  volume  of  a  gas 
may  be  referred.  The  pressure  universally  used  by 
chemists  for  this  purpose  is  the  pressure  of  a  column 
of  mercury  760  mm.  high  at  0°,  at  45°  latitude  and 
at  sea  level.  This  is,  approximately,  the  average 
pressure  of  the  air  at  sea  level  and  is  called  a  pres- 
sure of  one  atmosphere.  Other  pressures  are  most 
easily  determined  by  measuring,  directly  or  indi- 
rectly, the  height  of  the  column  of  mercury  which  will  balance 
the  elastic  pressure  of  the  gas. 

*  Corrections  for  Readings  of  the  Barometer.  —  In  accurate  work, 
when  the  barometer  is  read  at  some  other  temperature  than  0°  a  cor- 
rection must  be  subtracted,  owing  to  the  fact  that  the  column  of 
mercury  is  lighter  as  the  metal  expands  with  rise  of  temperature.  The 
correction  in  millimeters  at  temperatures  from  5°  to  35°  is  : 


Fig.  7 


TEMPEBATURE 

CORRECTION  FOR  BAROMETER 

CORRECTION  FOR  BAROMETER 

DEGREES 

WITH  GLASS  SCALE 

WITH  BRASS  SCALE 

5 

0.7 

0.6 

10 

1.3 

1.2 

15 

2.0 

1.9 

20 

2.6 

2.5 

25 

3.3 

3.1 

30 

4.0 

3.7 

35 

4.7 

4.3 

1  For  the  sake  of  simplicity,  the  lowering  of  the  meniscus  of  the 
mercury  in  the  tube  by  capillary  action  is  disregarded. 


LAWS  OF  GASES 


37 


If  the  pressure  is  less  than  760  mm.  the  correction  will  be  less  in  pro- 
portion. Thus  at  730  mm.  the  correction  for  a  glass  scale  is  2.5  mm.  at 
20°  instead  of  2.6  mm. 

The  corrections  for  latitude  and  altitude  are  usually  less  important. 

Corrections  of  barometer  for  latitude,  to  be  added  for  latitudes  less 
than  45°  or  subtracted  for  latitudes  greater  than  45°  : 


LATITUDE 

CORRECTION 

LATITUDE 

0° 

1.97 

90° 

5° 

1.94 

.85° 

10° 

1.85 

80° 

15° 

1.70 

75° 

20° 

1.51 

70° 

25° 

1.27 

65° 

30° 

0.98 

60° 

35° 

0.67 

55° 

40° 

0.34 

50° 

45° 

0.00 

45° 

Correction  for  altitude,  to  be  added. 


ALTITUDE 

CORRECTION 

BAROMETER  READING 

300  meters 

0.04 

720 

600  meters 

0.08 

700 

900  meters 

0.12 

680 

1200  meters 

0.16 

660 

1500  meters 

0.19 

640 

2000  meters 

0.24 

630 

If  the  corrections  for  latitude  and  altitude  are  applied  to  the  barometer 
readings,  the  weight  of  one  liter  of  the  gas  at  45°  latitude  may  be  properly 
used  in  calculating  the  weight  of  a  quantity  of  gas  measured  at  any  other 
latitude  or  altitude. 

A  problem  which  often  presents  itself  in  dealing  with  gases 
is  the  calculation  of  the  volume  which  a  quantity  of  gas,  that 
has  been  measured  at  some  other  pressure  than  that  of  one  at- 
mosphere, would  assume  if  it  were  brought  to  atmospheric  pres- 


38  A  TEXTBOOK  OF  CHEMISTRY 

sure.     Such  problems  are  most  easily  solved  by  putting  the 
mathematical  expression  given  above  into  the  following  form  : 

VP'  Pf 

V  =-^—~  °r  V  at  760  mm.  =  V  ^~ 

The  student  is  advised  most  earnestly  that  this  formula  should 
not  be  committed  to  memory.  Instead  of  this  it  should  only 
be  remembered  that  when  the  volume  at  one  pressure  is  given  and 
the  volume  at  another  pressure  is  desired,  the  first  volume  is  to 
be  multiplied  by  a  fraction  in  which  one  pressure  is  the  numera- 
tor and  the  other  pressure  the  denominator.  A  consideration  of 
the  fact  that  an  increase  in  pressure  will  cause  a  decrease  in  the 
volume  will  at  once  indicate  which  pressure  is  to  be  taken  as  the 
numerator  of  the  fraction.  The  proper  method  of  using  the 
formula  is  emphasized  because  in  the  study  of  chemistry  it  is 
of  the  greatest  importance  to  cultivate  the  ability  to  reason 
quickly  from  one  point  to  another  and  to  acquire  a  knowledge 
of  the  subject  by  a  rational  process  rather  than  by  mere 
memory. 

Effect  of  Temperature  on  a  Gas.  Law  of  Charles.  When 
the  temperature  of  a  gas  is  increased  one  degree  while  the  pres- 
sure remains  constant,  the  volume  will  increase  ^^  (or  0.003663) 
of  its  volume  at  O0.1  This  will  be  most  easily  understood  with 
the  aid  of  the  accompanying  diagram,  which  gives  the  volume 

xThis  law,  while  sufficiently  accurate  for  ordinary  calculations,  is 
only  approximate,  the  deviations  from  it  being  of  somewhat  the  same 
order  of  magnitude  as  the  deviations  from  the  law  of  Boyle.  The 
coefficients  of  expansion  of  some  of  the  more  common  gases  as  deter- 
mined by  the  increase  of  pressure  at  constant  volume  are  : 

Air  0.003666  or  1/272.8     Argon  0.003668  or  1/272.6 

Oxygen         0.003674  or  1/272.2     Helium  0.003663  or  1/273.0 

Nitrogen       0.003668  or  1/272.6     Carbon      mon- 
Nitric  oxide  0.003676  or  1/272.0         oxide  0.003667  or  1/272.7 

Hydrogen     0.003663  or  1/273.0     Carbon  dioxide  0.003698  or  1/270.4 

Sulfur  dioxide     0.003845  or  1/260.1 

As  with  the  law  of  Boyle,  those  gases  which  are  easily  liquefied  vary 
most  from  the  rate  of  expansion  for  a  "  perfect "  gas. 


ABSOLUTE  TEMPERATURES 


39 


373C 


283° 
273° 


173C 


73C 


Oc 


TEMPER- 
ATURE 

100° 


10°  --283cc. 


Oc 


-100°  --  173  cc. 

-200°  -\-    73  cc. 

-273° 
Fig.  8 


VOLUME 

373  cc. 


273  ec. 


which  273  cubic  centimeters  of  a  gas  at  0°  would  assume  at 
other  temperatures.  Only  hydrogen  or  helium  would  obey 
the  law  at  atmospheric  pressure  over  the  range  of  volumes 
given  in  the  diagram.  All  other  gases  are  liquid  or  solid 
at  -  200°. 

Absolute  Temperatures.  On 
the  left  side  of  the  diagram  is 
given  a  series  of  numbers  which 
are  called  absolute  temperatures. 
A  little  examination  of  the  dia- 
gram will  show  that  these  tem- 
peratures are  based  on  the 
thought  that  if  we  could  find 
a  gas  which  does  not  liquefy 
and  which  continued  to  con- 
tract at  the  same  rate  at  very 
low  temperatures  it  would  dis- 
appear at  -273°.  If  we  take 
this  point  as  the  starting  point 
for  the  "absolute"  scale  of 
temperature  it  is  evident  that 
the  freezing  point  of  water 
will  be  at  273°  absolute  and 
the  boiling  point  373°.  Any  temperature  may  be  readily 
converted  to  the  absolute  scale  by  adding  to  it,  algebraically  v 
273°. 

The  absolute  scale  of  temperature  enables  us  to  give  a  very 
simple  statement  of  the  law  of  Charles,  viz. :  The  volume  of  a 
gas  varies  directly  as  the  absolute  temperature.  This  becomes, 
mathematically : 

V :  V : :  T :  T',  or  F  =  — j- 

If,  as  is  customary  in  dealing  with  gases,  we  wish  to  find  the 
volume  which  a  gas,  which  has  been  measured  at  some  other 
temperature  than  0°,  would  assume  if  cooled  or  warmed  to  zero, 


40  A  TEXTBOOK  OF  CHEMISTRY 

273 
the  formula  may  be  written,  VQ=  V'——.     This  formula  should 

be  used  rationally,  not  by  rote  (see  p.  38),  and  may  be  combined 
with  the  formula  for  pressures  for  practical  uses.  If  the  volume 
of  a  gas  is  known  at  one  temperature  and  pressure,  its  volume  at 
some  other  temperature  and  pressure  may  be  calculated  by  multi- 
plying by  two  fractions  one  of  which  involves  the  two  pressures 
and  the  other  the  two  absolute  temperatures. 

Significance  of  the  Absolute  Zero.  The  absolute  scale  of 
temperature  may  be  treated  merely  as  a  mathematical  conven- 
ience in  dealing  with  problems  of  gases  and  of  thermodynamics ; 
but  the  question  naturally  arises  whether  the  absolute  zero  has 
any  further,  real  meaning.  Is  it,  in  reality,  as  the  name  indi- 
cates, a  point  of  absolute  cold  at  which  all  phenomena  of  tempera- 
ture begin  and  below  which  it  is  impossible  to  go  ?  Many  differ- 
ent phenomena  seem  to  indicate  that  the  absolute  zero  is  actual 
and  not  merely  a  mathematical  fiction.  It  can  be  no  mere  acci- 
dent that  hundreds  of  gases  and  vapors  obey  the  law  of  Charles  so 
closely ;  and  the  further  the  study  of  the  physical  and  chemical 
properties  of  gases  is  carried,  the  more  clear  does  it  become  that 
the  law  is  intimately  connected  with  some  of  the  most  funda- 
mental properties  of  matter.  From  the  side  of  experiment,  also, 
every  recent  attempt  to  reach  very  low  temperatures  has  pointed 
to  —  273°  as  a  point  which  can  never  be  passed.  The  lowest 
point  thus  far  reached  is  that  of  helium  boiling  under  a  pressure 
of  10  millimeters  and  is  estimated  as  —  270°,  or  3°  absolute. 
(Kamerlingh  Onnes,  Chemical  Abstracts,  1908,  p.  2752.) 

Determination  of  the  Weight  of  a  Liter  of  a  Gas.  The  weight 
of  a  unit  volume  of  any  gas  under  standard  conditions  is  one 
of  its  most  important  properties,  not  only  for  the  purpose  of  cal- 
culating the  weight  of  a  gas  when  we  know  its  volume,  but  be- 
cause of  relations  between  these  weights  for  different  gases,  upon 
which  one  of  the  most  important  laws  of  chemistry  is  based 
(p.  89).  The  unit  volume  usually  chosen  is  the  liter  and  the 
standard  conditions  are  a  temperature  of  zero  and  a  pressure  of 
760  mm.  of  mercury. 


WEIGHT  OF  GASES 


41 


If  a  bulb  1  is  connected  with  a  manometer  and  evacuated  by 
means  of  an  air  pump,  by  reading  the  manometer  and  tempera- 
ture, we  can,  if  we  know  the  capacity  of  the  bulb,  calculate  the 
volume  which  the  air  remaining  in  the  bulb  would  fill  at  0°  and 
760  mm.  If  we  weigh  the  bulb  and  then  fill  it  with  some  gas  at 


To  oirpurrfi 


Fig.  9 


atmospheric  pressure  (to  be  determined  by  reading  the  barom- 
eter) and  weigh  it  again,  the  difference  between  the  two 
weights  will  evidently  be  the  weight  of  the  gas  which  has  entered, 
while  the  volume  of  the  air  which  was  left  in  the  bulb  plus  that 
of  the  gas  which  has  entered  can  be  readily  calculated  for  stand- 
ard conditions  as  before.2  The  difference  between  this  calcu- 
lated volume  and  the  corrected  volume  of  the  air  which  re- 
mained in  the  bulb  will  be  the  volume,  under  standard  con- 
ditions, of  the  gas  which  was  admitted.  From  this  and  the 
weight  it  is  easy  to  calculate  the  weight  of  one  liter  of  the  gas 
under  standard  conditions. 

1  A  capacity  of  125  to  150  cc.  is  suitable  for  a  lecture  or  labora- 
tory experiment.      The  volume  may  be    determined  by  weighing 
the  bulb  empty  and  then  filled  with  water,  but  the  bulb  must  be 
thoroughly  dried  by  warming  it  and  evacuating  it  repeatedly  before 
it  is  used  for  the  determination. 

2  This  assumes,  of  course,  Dalton's  law  of  partial  pressures,  that 
when  two  gases  which  do  not  act  on  each  other  are  mixed,  each 
exerts  the  same  pressure  as  if  it  filled  the  whole  space  alone,  and  the 
total  pressure  is  the  sum  of  the  pressures  exerted  by  each  gas. 


42 


A  TEXTBOOK  OF   CHEMISTRY 


Graphical  Representation  of  the  Gas  Laws.  It  is  often  useful 
in  studying  physical  and  chemical  phenomena  to  use  a  method 
of  graphical  representation  which  is  illustrated  in  Figs.  10  and 
11.  In  Fig.  10  distances  from  the  line  OX  represent  pressures, 
while  distances  from  the  line  OY  represent  volumes.  If  we 


X 
V 

5 
4 

1s 
P 

2 
1 

( 

Jb 

i 

\ 

\ 

\ 

\ 

a 

\ 

^ 

^-^. 

"•^^ 

o 

r 

' 

)                   1                    2                    3                     4                    5PY 

PRESSURES 

Fig.  10 

start  with  a  unit  volume  of  a  gas  under  unit  pressure,  represented 
by  the  point  a,  as  the  pressure  increases  the  volume  will  decrease 
along  the  line  ac,  while  as  the  pressure  decreases  the  volume  will 
increase  along  the  lines  ab  and  PV,  the  product  of  pressure  and 
volume  must  always  remain  constant.  The  geometrical 
curve  which  satisfies  these  conditions  is  a  hyperbola. 

Charles's  law  may  be  represented  in  a  similar  way  by  Fig.  11. 


LAWS  OF  GASES 


43 


Here  the  relation  between  volumes  and  absolute  temperatures 
is  represented  by  a  straight  line,  but  all  gases  liquefy  before  the 


500 


400 


£300 

200 


100 


V 


O 


100 c 


200°  300°  400 

ABSOLUTE  TEMPERATURES 


500° 


Fig.  11 


absolute  zero  is  reached,  and  the  line  can  never  be  continued, 
experimentally,  to  the  origin.1    . 


EXERCISES 

1.  A  quantity  of  gas  fills  a  volume  of  175  cc.  at  20°  and  under  a 
pressure  of  735  mm.      What  will  be  its  volume  under  standard  con- 
ditions (0°  and  760  mm.)  ? 

2.  A  flask  having  a  capacity  of  3.5  liters  is  filled  with  oxygen  at 
25°  and  770  mm.     What   weight   of   oxygen  does   it  contain  ?     (See 
p.  22.) 

3.  What  volume  will  33.  5  cc.  of  a  gas  measured  at  18°  and  715  mm. 
assume  at  25°  and  731  mm.  ? 

4.  A  cylindrical  gasometer  has  a  diameter  of  30  cm.  and  height  of 
60  cm.     What  weight  of  oxygen  will  be  required  to  fill  it  at  22°  and 
745  mm.  ? 

5.  A  bulb  having  a  capacity  of  127.2  cc.  was  exhausted  till  the 
manometer  showed  a  pressure  of  35  mm.  while  the  temperature  was 

1  The  line  0  Y  is  called  the  axis  of  abscissas  and  any  line  parallel  to 
it  and  perpendicular  to  OX  is  called  an  abscissa,  while  OX  is  the  axis 
of  ordinates  and  any  line  parallel  to  this  and  perpendicular  to  Or 
is  called  an  ordinate.  O  is  called  the  origin. 


44  A  TEXTBOOK  OF  CHEMISTRY 

23°.  After  weighing,  it  was  filled  at  atmospheric  pressure  with  a  gas. 
The  reading  of  the  barometer  was  751  mm.  and  the  temperature  23°, 
as  before.  The  increase  in  weight  was  0.1382  gram.  What  is  the  weight 
under  standard  conditions  of  one  liter  of  the  gas  which  was  used  ? 

(Ans.    1.2507.) 

6.  A  sample  of  bituminous  coal  has  the  following  composition : 
Carbon  75.00  per  cent 

Hydrogen  5.25  per  cent 

Oxygen  10.00  per  cent 

Ash,  nitrogen,  etc.        9.75  per  cent 
100.00  per  cent 

What  is  the  heat  of  combustion  of  the  coal  in  calories  per  kilo- 
gram and  in  B.  T.  U.  per  pound,  assuming  the  heat  of  combustion 
of  one  gram  of  carbon  as  8080  calories,  one  gram  of  hydrogen  as 
34,179  calories  (burned  to  liquid  water),  and  that  1.25  per  cent  of  the 
hydrogen  is  combined  with  the  oxygen  and  contributes  nothing  to 
the  heat  of  combustion  ?  It  is  assumed  further  that  the  combination 
between  the  carbon  and  hydrogen  is  of  such  a  nature  that  these  ele- 
ments give  the  same  amount  of  heat  when  the  coal  is  burned  as  they 
would  give  if  they  were  in  the  free  state. 


CHAPTER  IV 
HYDROGEN 

SYMBOL,  H.    ATOMIC  WEIGHT,  1.0078. 

Occurrence.  Although  the  quantity  of  hydrogen  in  the  world 
is  very  much  smaller  than  the  quantity  of  oxygen,  it  is  very 
widely  diffused,  especially  in  the  form  of  its  most  common  com- 
pound, water.  It  forms  a  little  more  than  one  ninth  of  the 
weight  of  water  and  is  present  both  as  water  and  as  a  constitu- 
ent of  all  of  the  most  important  compounds  found  in  vegetables 
and  animals.  Hydrogen  is  an  essential  element,  also,  in  the 
large  class  of  substances  called  acids.  A  minute  quantity,  pos- 
sibly 0.001  per  cent,  or  one  part  in  100,000,  is  found  free  in  the 
air.  There  is  some  evidence  that  at  very  high  altitudes  the  at- 
mosphere consists  almost  exclusively  of  hydrogen. 

Acids.  In  order  to  understand  one  of  the  most  convenient 
methods  for  the  preparation  of  hydrogen  in  the  laboratory,  it 
is  necessary  to  know  something  of  the  properties  of  the  important 
class  of  substances  called  acids.  It  has  been  shown  that  when 
sulfur  is  burned  in  the  air,  sulfur  dioxide,  SO2,  is  formed.  A 
small  amount  of  the  sulfur  usually,  or  perhaps  always,  combines 
with  more  oxygen  to  form  sulfur  trioxide,  80s,  and  by  means  of 
suitable  apparatus  and  a  catalytic  agent,  nearly  all  of  the  sulfur 
can  be  converted  into  this  compound  (see  p.  175).  When  sulfur 
trioxide  is  dissolved  in  water  it  combines  with  it,  giving  sulfuric 
acid : 

SO3    +    H2O    =    H2S04 

-Sulfur  Water  Sulfuric 

Trioxide  Acid 

It  will  be  recalled  that  this  sort  of  combination  between  oxides 
of  nonmetallic  elements  and  water  gave  to  chemistry  the  name 

45 


46 


A  TEXTBOOK  OF  CHEMISTRY 


of  oxygen.  Sulfuric  acid,  when  pure,  is  a  heavy  liquid  of 
an  oily  consistency,  sometimes  called  oil  of  vitriol. 

By  the  action  of  sulfuric  acid  on  common  salt  we  can  obtain 
hydrochloric  acid,  HC1,  a  gas  which  dissolves  easily  in  water  and 
which  is  ordinarily  used  in  the  form  of  its  solution. 

By  the  action  of  sulfuric  acid  on  saltpeter,  nitric  acid,  HNOs, 
is  formed.  This  is  a  liquid,  which  is  usually  diluted  with  water 
for  use. 

The  most  common  acid  of  ordinary  experience  is  acetic  acid, 
HC2HaO2,  the  acid  of  vinegar.  This  is  the  acid  from  which  we 
have  all  learned  to  associate  the  word  acid  with  the  sour  taste 
which  is  characteristic  of  all  moderately  strong  acids. 

Radicals.  An  examination  of  the  formulas  of  the  acids 
named  above  shows  that  each  of  them  contains  hydrogen,  but  a 
still  more  important  characteristic  of  these  and  of  all  other  acids 
is  that  in  a  great  variety  of  reactions  this  hydrogen  may  be  re- 
placed by  other  elements  and  especially  by  metals.  The  follow- 
ing are  illustrations  of  such  replacement : 


H2SO4 

+         Zn 

ZnSO4        4 

2H 

Sulfuric 

Zinc 

Zinc  Sulfate 

Hydrogen 

Acid 

HC1 

+        Na 

NaCl         4 

H 

Hydrochloric 

Sodium 

Sodium  Chloride 

Acid 

(Common  Salt) 

HNO3 

+     NaOH 

=      NaN03 

h    H20 

Nitric  Acid 

Sodium 

Sodium 

Water 

Hydroxide 

Nitrate 

HC2H3O2 

+     NaOH 

=      NaC2H302     - 

h     HOH 

Acetic 

Sodium 

Sodium 

Water 

Acid 

Hydroxide 

Acetate 

In  each  of  these  reactions  one  or  two  atoms  of  hydrogen  are 
replaced  by  an  atom  1  of  some  metal,  while  all  of  the  rest  of  the 

1  One  of  the  most  common  mistakes  of  beginners  in  such  cases 
is  to  say  "one  or  two  parts  of  hydrogen  are  replaced  by  one  part 
of  the  metal."  The  distinction  between  one  part  and  one  atomic 
weight  (or  in  accordance  with  the  atomic  theory,  one  atom)  of  an 
element  ought  always  to  be  kept  clear. 


HYDROGEN  47 

acid  passes  into  the  new  compound  without  any  change  in  com- 
position. A  group  of  atoms,  which  remain  in  combination  in 
this  way  when  they  pass  from  one  compound  to  another,  is 
called  a  radical.  Thus  SO4  is  the  radical  of  sulfuric  acid ;  NO3, 
of  nitric  acid ;  C2H3O2,  of  acetic  acid. 

Salts.  The  compounds  formed  by  the  replacement  of  the 
hydrogen  of  an  acid  by  a  metal  are  called  salts.  These  are  so 
intimately  connected  with  the  acids  in  their  composition  that  it 
is  natural  to  use  for  them  names  which  are  derived  from  the 
names  of  the  acids.  How  this  is  done  is  clear  from  the  illustra- 
tions given.  The  name  of  the  metal  of  the  salt  is  placed  first 
and  this  is  followed  by  a  word  in  which  the  -ic  of  the  acid  is 
changed  to  -ate.  Sulfuric  acid  gives  sulfates;  nitric  acid,  nitrates ; 
acetic  acid,  acetates.  The  name  o£  common  salt,  sodium 
chloride,  seems  to  be  an  exception,  but  this  is  because,  as  a  binary 
compound,  it  belongs  to  the  class  of  substances  which  take 
names  ending  in  -ide  (p.  29).  Additional  principles  which  are 
used  in  naming  acids  and  salts  will  be  considered  later. 

Preparation  of  Hydrogen.  1.  Electrolysis  of  Dilute  Sulfuric 
Acid.  If  an  electrical  current  is  passed  between  two  strips  of 
platinum  (called  electrodes)  which  are  immersed  in  dilute  sul- 
furic acid,  bubbles  of  gas  will  rise  from  the  electrodes ;  and  if  an 
apparatus  is  so  arranged  (p.  9)  that  these  can  be  collected, 
it  will  be  found  that  the  gas  rising  from  the  negative  electrode 
(cathode)  is  hydrogen,  while  that  from  the  positive  electrode 
(anode)  is  oxygen.  The  volume  of  the  hydrogen  will  be  almost 
exactly  twice  that  of  the  oxygen. 

Since,  as  we  shall  find  later,  the  hydrogen  and  oxygen  are  lib- 
erated in  the  same  proportion  in  which  they  combine  to  form 
water,  this  experiment  is  often  spoken  of  as  a  decomposition 
of  water  by  electricity  —  and  in  a  certain  sense  this  is  correct. 
That  the  sulfuric  acid  is  more  than  a  merely  passive  agent  in 
what  takes  place  is  evident,  however,  first,  because  pure  water 
is  nearly  a  nonconductor  for  electricity ;  and,  second,  because  if 
we  examine  the  liquid  in  the  U-tube  by  appropriate  means,  we 
shall  find  that  the  hydrogen  atoms  of  the  sulfuric  acid  are  being 


48"  A  TEXTBOOK  OF  CHEMISTRY 

transferred  through  the  liquid  toward  the  cathode  as  the  current 
passes,  while  the  radicals  of  the  sulfuric  acid,  the  SO4  group  of 
atoms,  are  transferred  toward  the  anode.  In  other  words,  elec- 
trolysis is  not  merely  something  which  takes  place  at  the  two 
electrodes,  but  it  is  always  accompanied  by  a  transfer  of  material 
through  the  whole  of  the  space  between ;  and  while  hydrogen  is 
carried  in  one  direction,  it  is  the  sulfate  radical  and  not  oxygen, 
which  is  carried  the  other  way. 

Electrolytes.  Ions.  Theory  of  electrolysis.  Any  substance 
which  carries  the  electric  current  in  this  way  is  called  an  elec- 
trolyte. The  most  satisfactory  theory  which  has  been  proposed 
to  explain  the  facts  which  have  just  been  given  is  that  electro- 
lytes in  solution  are  more  or  less  completely  separated  into  parts 
which  are  charged  with  positive  or  negative  electricity.  Ac- 
cording to  this  theory  sulfuric  acid  separates  partly  into  hydro- 
gen atoms  with  a  positive  charge  of  electricity  and  the  sulfate 
radical  with  two  negative  charges.  This  is  indicated  by  the 
symbols  H+,  H+,  SO 4  .  When  the  positive  and  negative  elec- 
trodes are  dipped  in  the  dilute  acid,  the  positively  charged  hy- 
drogen atoms  are  attracted  by  the  negative  cathode  and  move 
toward  it,  while  the  negatively  charged  sulfate  radicals  are  re- 
pelled by  the  cathode  and  attracted  by  the  positive  anode. 
This  causes  the  motion  of  the  hydrogen  atoms  in  one  direction 
and  the  motion  of  the  sulfate  radicals  in  the  other,  through  the 
solution.  This  motion  constitutes  the  current  of  electricity  in 
an  electrolyte.  At  the  cathode  the  hydrogen  atoms  lose  their 
positive  charge  and  at  once  appear  as  hydrogen  gas.  At  the 
anode  the  action  is  more  complicated,  but  the  final  result  i& 
that  oxygen  gas  is  liberated.  The  charged  atoms  or  groups 
are  called  ions.  The  positive  ion  is  called  the  cation,  the  neg- 
ative ion,  the  anion.  The  decomposition  of  an  electrolyte  by 
an  electric  current  is  called  electrolysis. 

2.  Preparation  of  Hydrogen  from  Iron  and  Steam.  If  steam 
is  passed  over  red-hot  iron  contained  in  an  iron  tube  (Fig.  12), 
a  part  of  it  gives  up  its  oxygen  to  the  iron  and  hydrogen  is 
liberated.  If  the  compound  which  remains  in  the  tube  is  ex- 
amined, it  is  found  to  have  the  same  composition  as  the  magnetic 


HYDROGEN 


49 


oxide  of  iron,  Fe3O4,  formed  when  iron  burns  in  oxygen.     The 
equation  is  not  quite  so  simple  as  those  which  have  been  given 


Fig.  12 

before.  In  order  to  arrive  at  the  true  equation  the  formulas  of 
the  substances  used  and  the  products  obtained  should  be  written 
first : 

Fe      +       H2O      ->     Fe3O4      +       H 

Iron  Water  Magnetic  Hydrogen 

Oxide  of  Iron 

On  examining  the  above  it  is  seen  that  4  atoms  of  oxygen  will 
be  required  to  form  one  molecule  of  the  magnetic  oxide  of  iron, 
hence  we  must  have  4  molecules  of  water  in  the  first  member  of 
the  equation  to  furnish  these.  The  4  molecules  of  water  will 
give  8  atoms  of  hydrogen  and  3  atoms  of  iron  will  also  be  re- 
quired to  form  the  magnetic  oxide.  Putting  all  together  we 
have: 

3  Fe  +  4  H2O  =  Fe3O4  +  8  H 

It  would,  doubtless,  be  easier  for  a  beginner  to  learn  this  last 
equation  outright  than  to  learn  how  to  derive  it  in  the  manner 
indicated,  but  things  which  are  a  mere  matter  of  memory  are 
likely  to  be  evanescent,  while  a  rational  process  like  the  above 
can  be  reproduced  at  will.  It  is  very  important  in  studying 
chemistry  to  reduce  those  portions  which  are  remembered  as 
distinguished  from  those  portions  which  are  derived  by  a  logical 


50  A  TEXTBOOK  OF  CHEMISTRY 

process  just  as  far  as  possible.  At  the  same  time  many  simple, 
fundamental  facts,  as  here  the  composition  of  the  magnetic 
oxide  of  iron,  must  be  remembered  and  used  over  and  over  again. 

Reversible  Reactions.  It  was  stated  above  that  a  part  only 
of  the  steam  is  decomposed  by  the  iron.  If  we  reverse  the  con- 
ditions and  pass  hydrogen  over  magnetic  oxide  of  iron,  part  of 
the  hydrogen  will  be  converted  into  water  and  metallic  iron  will 
be  obtained.  The  earlier  and  most  natural  idea  of  chemical 
affinity  was  that  when  three  elements  are  present  those  two 
would  unite  which  had  the  strongest  affinity  for  each  other. 
If  this  were  true,  either  the  hydrogen  would  be  able  to  take  the 
oxygen  from  the  iron  or  the  iron  should  be  able  to  take  it  from  the 
hydrogen.  We  see  from  the  experiments  described  that  this 
simple  idea  is  not  correct,  but  that  either  element  can  take  the 
oxygen  from  the  other.  While  there  is  a  certain  sense  in  which 
iron  has  a  stronger  affinity  for  oxygen  than  hydrogen  has,  the 
direction  of  the  reaction  depends  on  the  quantities  of  the  sub- 
stances present  as  well  as  upon  their  relative  affinities.  If 
steam  is  used  and  the  hydrogen  is  constantly  removed,  the  tend- 
ency is  to  form  magnetic  oxide  of  iron  and  hydrogen.  If  hydro- 
gen is  used  and  the  steam  is  constantly  removed,  the  tendency 
is  to  form  metallic  iron  and  water. 

It  is  often  convenient  to  express  such  reversible  reactions  as 
follows  :  3  Fe  +  4  H2O  ^±  Fe3O4  +  8  H 

3.  Decomposition  of  Water  by  Metals  at  Ordinary  Tempera- 
tures. Potassium  and  sodium  have  a  much  stronger  affinity 
for  oxygen  than  iron  has,  and  partly  for  this  reason,  partly, 
perhaps,  for  other  reasons  which  we  do  not  fully  understand, 
these  metals  will  decompose  water  and  liberate  hydrogen  at 
ordinary  temperatures.  If  potassium  is  thrown  on  water,  the 
heat  of  the  reaction  is  great  enough  to  cause  the  hydrogen  to 
ignite.  It  burns  with  a  violet  flame,  the  color  being  given  to  it 
by  the  potassium.  Sodium  when  thrown  on  water  usually  rolls 
over  the  surface  in  a  globule,  evolving  hydrogen,  which  does  not 
take  fire,  but  if  thrown  on  a  piece  of  filter  paper  lying  on  the 


HYDROGEN 


51 


water  so  that  the  globule  remains  at  one  spot,  the  hydrogen  will 
catch  fire  and  burn  with  the  yellow  flame  characteristic  of  sodium. 
If  a  small  piece  of  sodium  is  wrapped  in  paper  and  thrust  quickly 
under  the  mouth  of  a  jar  which  has  been 
filled  with  water  and  inverted  with  the 
mouth  under  water,  the  sodium  will  act  on 
the  water  as  before  and  the  hydrogen  may 
be  collected  and  examined. 

If  the  water  in  which  the  potassium  or 
sodium  has  dissolved  in  these  experiments  is 
examined,  it  will  be  found  to  have  a  soapy 
feel  and  disagreeable,  acrid  taste.  It  will 
also  turn  the  color  of  red  litmus  paper  blue. 
If  the  solution  is  evaporated  in  a  dish  of 
platinum  or  of  some  material  which  is  not 
affected  by  it  and  under  such  conditions  that  it  cannot  absorb 
carbon  dioxide  from  the  air,  a  white  solid  will  be  obtained, 
which  will  have  the  composition  represented  by  the  formula 
KOH  or  NaOH.  These  substances  absorb  and  retain  water 
so  greedily  that  it  is  necessary  to  heat  them  nearly  to  redness 
before  the  last  of  the  water  is  expelled.  They  are  called,  in 
accordance  with  their  composition,  potassium  hydroxide  or 
sodium  hydroxide.  The  equations  are  : 

K          +    H2O      =      KOH       +        H 

Potassium  Water  Potassium  Hydrogen 

Hydroxide 


Fig.  13 


Na        +    H20      = 
Sodium 


NaOH      + 

Sodium 
Hydroxide 


H 


or, 


Na        +    HOH    =      NaOH 

Hydrogen 
Hydroxide 


+ 


Contrast  between  the  Action  of  Iron  and  of  Sodium  on  Water. 
The  last  form  expresses  a  little  more  clearly  that  the  metal  has 
replaced  only  one  of  the  two  atoms  of  hydrogen  in  each  molecule 


52  A  TEXTBOOK  OF  CHEMISTRY 

of  water.  The  action  is  seen  to  be  quite  different  from  that  of 
iron  on  steam.  This  is  closely  connected  with  the  amount  of 
chemical  energy  changed  to  heat  in  each  reaction.  In  the  reac- 
tion, Na  +  H2O  =  NaOH  +  H,  43,450  calories  are  liberated 
for  each  gram  atom  of  hydrogen  set  free,  while  in  the  reaction, 
3  Fe  +  4  H2O  =  Fe3O4  +  8  H,  if  it  could  be  carried  out  at  100°, 
only  4160  calories  would  be  given  for  each  gram  atom  of  hydro- 
gen liberated.  In  general,  those  reactions  in  which  large 
amounts  of  chemical  energy  are  changed  to  heat  take  place  most 
easily.  We  must,  however,  guard  against  the  impression  that 
this  is  a  universal  law.  The  ease  with  which  a  reaction  takes 
place  is  by  no  means  proportional  to  the  heat  generated.  Other 
factors  are  involved,  and  some  of  these  are,  at  present,  but  little 
understood. 

4.  Hydrogen  from  "  Hydrone."     The  action  of  water  on  so- 
dium is  too  violent  for  use  as  a  laboratory  method  of  preparing 
hydrogen  in  quantity.     If,  however,  the  sodium  is  alloyed  with 
lead,  the  action  is  moderated,  and  such  an  alloy  containing  about 
35  per  cent  of  sodium  is  sold  under  the  name  of  "  hydrone." 
By  means  of  it  very  pure  hydrogen  can  be  easily  prepared. 

5.  Preparation  of  Hydrogen  by  the  Action  of  Metals  on  Acids. 
If  a  strip  of  zinc  and  one  of  platinum,  copper  or  lead  are  dipped 
in  dilute  sulfuric  or  hydrochloric  acid  while  the  strips  are  con- 
nected by  means  of  a  wire,  an  electrical  current  will  pass  through 
the  wire  while  bubbles  of  hydrogen  will  be  seen  to  collect  and  rise 
from  the  surface  of  the  platinum,  copper  or  lead.     If  the  liquid 
between  the  two  metallic  plates  is  examined,  as  in  the  electrolysis 
of  dilute  sulfuric  acid,  it  will  be  found  that  the  hydrogen  travels 
through  the  liquid  toward  the  platinum,  while  the  sulfate  radical 
or  the  chlorine  travels  toward  the  zinc.     At  the  surface  of  the 
zinc,  the  sulfate  radical  or  the  chlorine  combines  with  the  zinc, 
forming  zinc  sulfate,  ZnSO4,  or  zinc  chloride,  ZnC^.     If  pieces 
of  chemically  pure  zinc  are  placed  in  dilute  hydrochloric  or  sul- 
furic acid,  there  will  be  almost  no  action  at  all,  while  commercial 
zinc  will  dissolve  rapidly.     After  the  action  of  the  acid  on  the 
commercial  zinc  has  continued  for  a  short  time  it  will  be  seen 


HYDROGEN  53 

that  the  surface  is  dark,  and  a  closer  examination  will  show  that 
it  is  covered  with  lead  and  other  impurities  found  in  the  zinc. 
When  we  consider  these  facts  along  with  the  experiment  with  the 
strips  of  platinum  and  zinc,  we  reach  the  conclusion  that  the  ac- 
tion of  the  acid  on  the  zinc  requires  some  catalytic  agent  like 
the  platinum  or  lead  before  it  can  be  very  rapid,  and  that  the 
phenomenon  of  the  solution  of  the  zinc  is  partly  electrical,  being 
accompanied  by  electrical  currents  between  the  particles  of  lead 
and  zinc  in  the  impure  zinc.  If  we  disregard  the  catalytic  agent, 
the  process  may  be  expressed  by  the  equations : 

Zn    +     H2SO4        =       ZnSO4     +       2H 

Zinc  Sulfuric  Zinc  Hydrogen 

Acid  Sulfate 

Zn    +     2HC1        =       ZnCl2      +       2H 

Hydrochloric  Zinc 

Acid  Chloride 

It  is  to  be  noticed  that  one  atom  of  zinc  replaces  two  atoms  of 
hydrogen  in  each  case  and  that  when  an  acid  is  used  which  has 
only  one  atom  of  hydrogen  in  its  molecule,  two  molecules  of  the 
acid  are  required  for  the  reaction. 

If  iron  is  substituted  for  zinc,  these  reactions  become : 


Fe    + 

H2SO4 

=     FeSO4 

+      2H 

Iron 

Sulfuric 

Ferrous 

Hydrogen 

Acid 

Sulfate 

Fe    + 

2HC1 

=     FeCl2 

+      2H 

Ferrous 

Chloride 

Apparatus  for  the  Preparation  of  Hydrogen.  In  the  labora- 
tory, small  quantities  of  hydrogen  may  be  generated  in  the 
simple  apparatus  shown  in  Fig.  14.  Zinc  and  some  water  are 
placed  in  the  generating  flask,  and  dilute  sulfuric  acid1  is  added 
in  portions  through  the  thistle  tube. 

1  As  sulfuric  acid  is  heavier  than  water  (sp.  gr.  1.84)  and  much 
heat  is  generated  on  its  dilution,  it  should  always  be  poured  slowly 
into  water  and  should  never  be  diluted  by  pouring  water  upon  the 
acid.  Pouring  water  on  concentrated  sulfuric  acid  may  cause  an 
explosion. 


54 


A  TEXTBOOK  OF  CHEMISTRY 


A  more  convenient  apparatus  for  the  preparation  of  larger 
amounts,  or  when  it  is  desired  to  have  the  gas  always  ready  for 
use,  is  the  Kipp  generator  (Fig.  15).  The 
zinc  is  placed  in  the  middle  bulb  and  the 
dilute  acid  is  poured  in  through  the  upper 
bulb,  which  communicates  with  the  lower  one 
through  the  tube  A.  When  the  stopcock  B 
is  opened,  the  acid  rises  and  comes  in  contact 
with  the  zinc  in  the  middle  bulb  and  the  gen- 
eration of  hydrogen  begins.  Whenever  the 
stopcock  is  closed  the  hydrogen  generated 
forces  the  acid  away  from  the  zinc  and  the 
action  ceases  as  soon  as  the  acid  moistening 
the  surface  of  the  zinc  is  exhausted.  The 


Fig.  14 


generator  is  not  altogether  satisfactory  because  the  spent  acid 
containing  zinc  sulfate  is  mixed  with  that  which  has  not  been 
used,  diluting  it  and  causing  the 
action  to  become  very  slow  before 
the  acid  has  been  exhausted.  A 
more  suitable  form  of  apparatus 
for  generating  large  quantities  of 
hydrogen  is  described  on  p.  165. 

Purification  of  Hydrogen.  The 
hydrogen  prepared  by  any  of  the 
methods  described  is  impure. 
Spray  from  the  generating  liquids 
may  be  removed  by  passing  the 
gas  through  a  tube  filled  with  cot- 
ton wool.  Moisture,  or  water 
vapor  may  be  removed  by  means 
of  calcium  chloride,1  contained  in  a 


Fig.  15 


1  One  liter  of  a  gas  dried  with  cal- 
cium chloride  retains  1.0  mg.  of  water 
at  15°,  1.5  mg.  at  20°,  2.5  mg.  at. 

25°,  3.3  mg.  at  30°.  When  dried  with  concentrated  sulfunc  acid 
the  amount  of  water  retained  by  one  liter  of  the  gas  is  only 
0.002  mg.  at  15°  to  19°.  When  dried  with  phosphorus  pentoxide 


HYDROGEN 


55 


tube  such  as  shown  in  Fig.  16,  or  more  perfectly  by  means  of 
pumice  stone  or  glass  beads  moistened  with  concentrated  sul- 
furic  acid  or  by  phosphorus  pentoxide.  Hydrogen 
sulfide  1  and  some  other  impurities,  especially  some 
of  those  which  give  an  unpleasant  odor  to  the  gas, 
may  be  removed  by  passing  it  through  a  wash  bottle 
containing  a  solution  of  potassium  permanganate, 
but  a  small  amount  of  oxygen  will  be  introduced 
into  the  gas  (V.  Meyer,  and  Recklinghausen  Ber.  29, 
2550).  Oxygen  may  be  removed  by  passing  the  gas 
through  a  hot  tube  containing  platinized  quartz, 
which  will  cause  the  oxygen  to  combine  with  some  of 
the  hydrogen.  This  should  be  done,  of  course,  be- 
fore the  gas  is  dried.  Nitrogen 
from  the  air  cannot  be  removed 
and  when  pure  hydrogen  is  re- 
quired very  great  care  is  necessary 
to  prevent  its  entrance  (Cooke  ji-  jg 
and  Richards,  Am.  Chem.  J.  10, 
81;  Morley,  ibid.  17,  267;  Noyes,  J.  Am. 
Chem.  Soc.  30,  1724). 

Properties  of  Hydrogen.  Hydrogen  is  a 
colorless,  tasteless  and  odorless  gas.  One 
liter  weighs  at  0°  and  760  mm.  pressure 
0.08987  gram.  As  the  weight  of  a  liter  of 

air  is  1.293  grams,   air  is 


Fig.  17 


0.08987 
times  heavier  than  hydrogen,  or  approxi- 


mately 14 J  times.     Oxygen  is  — ! =  15.90    times    heavier 

than  hydrogen,  or  approximately  16  times. 


the  amount  of  moisture  retained  by  one  liter  of  the  gas  is  less  than 
0.00002  mg.  (Morley). 

1  Hydrogen  sulfide  may  be  removed  to  better  advantage  by 
passing  the  gas  through  a  wash  bottle  or  serpentine  tube  containing 
lead  oxide  dissolved  in  a  solution  of  potassium  hydroxide. 


56 


A  TEXTBOOK  OF  CHEMISTRY 


Fig.  18 


At  a  very  low  temperature  hydrogen  condenses  to  a  liquid 
which  boils  under  atmospheric  pressure  at  —  252.5°  or  20.5°,  ab- 
solute. The  liquid  has  a  density  of  only  0.07  gram  per  cubic 
centimeter  or  one  gram  fills  a  volume  of  about  14  cc.  If  the 

liquid  is  made  to  boil  by  reduc- 
ing the  pressure,  it  grows  still 
colder  and  at  -  260°,  or  13°,  ab- 
solute, what  remains  freezes  to 
a  solid.  The  vapor  pressure  of  the 
solid  is  58  mm. 

The  lightness  of  hydrogen  as 
compared  with  air  may  be  easily 
shown  by  pouring  it  upward 
through  the  air,  by  showing  that 
it  may  be  collected  in  an  inverted 
jar  or  beaker  while  it  will  not  remain  in  one  which  is  placed 
upright,  and  by  filling  soap  bubbles  or  toy  balloons  with  the  gas. 
The  use  of  the  gas  for  filling  balloons  is  well  known. 
*  What  is  the  lifting  power  of  one  cubic  meter 
of  pure  hydrogen  under  standard  conditions? 
What  will  be  the  lifting  power  of  one  cubic  meter 
at  a  mile  above  sea  level  when  the  pressure 
is  620  mm.  and  the  temperature  20°  ?  It  is  con- 
venient to  remember  that  as  the  volume  of  a  gas 
varies  inversely  with  the  pressure  and  directly 
with  the  absolute  temperature,  the  weight  of  a 
given  volume  must  vary  directly  as  the  pressure 
and  inversely  as  the  absolute  temperature. 

Diffusion  of  Gases.  If  two  cylinders  are  filled, 
one  with  hydrogen  and  the  other  with  air,  and 
placed  with  their  mouths  together  but  with  the 
hydrogen  above  (Fig.  19),  it  will  be  found  after  a 
comparatively  short  time  that  the  gas  in  each 
cylinder  will  explode,  if  ignited,  with  the  whistling  sound  char- 
acteristic of  a  mixture  of  air  and  hydrogen  fired  in  an  open 
cylinder  or  test  tube.  Although  the  air  is  fourteen  and  a  half 


Fig 


DIFFUSION 


57 


times  heavier  than  the  hydrogen,  it  makes  its  way  quite  rapidly 
up  into  the  hydrogen  above  it  and  the  hydrogen  passes  down 
into  the  heavier  air  below.  This  property  of  mixing  with  each 
other  is  true  of  all  gases  without  exception  and  is  called 
diffusion.  While  there  are  many  liquids  which  do  not  dissolve 
in  each  other  or  which  dissolve  only  to  a  limited  extent  and 
such  liquids  separate  into  layers  in  accordance  with  their 
specific  gravities,  gases,  which  differ  much  more  than  liquids  in 
their  densities,  always  mix  when  brought 
in  contact  and  when  once  mixed  will  sep- 
arate to  only  a  very  slight  extent  (p.  45) 
in  accordance  with  their  densities.  Very 
closely  related  to  the  diffusion  of  gases 
which  are  in  contact  with  each  other  is 
the  diffusion  of  gases  through  a  wall  full 
of  fine  openings,  which  separates  them. 
If  a  cylinder  of  porous  porcelain  with 
openings  so  fine  that  pressure  will  cause  a 
gas  to  pass  through  them  only  very  slowly 
is  fitted  with  a  rubber  stopper  and  con- 
nected with  a  bulb  and  bent  tube  filled 
with  water  as  shown  in  Fig.  20,  on  bring- 
ing a  beaker  filled  with  hydrogen  over 
the  cylinder  the  pressure  within  will  sud- 
denly increase  and  force  water  out  of  the 
tube  in  a  jet,  showing  that  hydrogen  is 


Fig.  20 


passing  through  the  walls  of  the  cylinder  to  the  interior. 
When  the  stream  of  water  ceases,  if  the  beaker  is  removed, 
the  movement  of  the  water  in  the  tube  will  show  that  diffu- 
sion of  the  hydrogen  outward  is  taking  place.  By  appro- 
priate experiments  it  is  possible  to  show  that  some  air  passes 
out  through  the  porous  wall  while  the  hydrogen  is  passing  in. 
A  careful  study  of  the  phenomena  by  Graham  has  shown  that 
gases  pass  through  a  porous  wall  of  this  sort,  at  a  rate  which 
varies  inversely  as  the  square  root  of  the  density.  Oxygen, 
which  is  about  sixteen  times  as  heavy  as  hydrogen,  will  pass 
the  wall  only  one  fourth  as  fast. 


58  A  TEXTBOOK   OF  CHEMISTRY 

Kinetic  Theory  of  Gases.  When  water  is  converted  into 
steam  at  100°  the  volume  at  atmospheric  pressure  is  increased 
more  than  1600  times.  This  fact  and  many  of  the  other  proper- 
ties of  gases  makes  it  seem  highly  probable  that  the  space  be- 
tween the  molecules  of  a  gas  is  very  large  in  comparison  with  the 
size  of  the  molecules  themselves.  A  very  satisfactory  explana- 
tion of  the  law  of  diffusion  of  gases,  given  in  the  last  paragraph, 
and  of  the  fact  that  a  gas  expands  at  once  to  fill  any  empty  space, 
however  large,  which  it  is  allowed  to  enter,  is  found  in  the  kinetic 
theory  of  gases.  According  to  this  theory  the  molecules  of 
gases  are  moving  constantly  at  a  comparatively  high  velocity, 
and  whenever  they  meet  each  other  they  rebound  according  to 
the  laws  of  elastic  bodies.  According  to  these  laws  when  two 
elastic  bodies  meet  each  other  squarely  each  rebounds  with  the 
energy  of  the  other.  If  the  bodies  meet  at  an  angle,  the  inter- 
change of  energies  will  be  only  partial,  and  the  effect  of  this 
constant  interchange  must  be  to  give  to  all  molecules  of  the 
same  weight  approximately  the  same  average  energy  and  hence 
the  same  average  velocity.  If  molecules  of  different  weights 
are  mixed,  however,  the  interchange  of  energies  must  give  a 
greater  velocity  to  lighter  molecules.  As  the  energy  of  a  moving 
body  varies  as  the  square  of  its  velocity,  the  velocities  of  mole- 
cules of  different  weights  must  vary  inversely  as  the  square 
roots  of  their  weights,  if  their  energies  are  the  same,  since  the 
pressure  of  a  gas  is  due  to  the  impacts  of  its  molecules  on 
the  walls  of  the  containing  vessel  and  the  pressure  does  not 
change  when  two  gases  having  different  densities  are  mixed. 

As  an  illustration  we  may  take  oxygen  and  hydrogen.  The 
molecule  of  oxygen  is  16  times  as  heavy  as  that  of  hydrogen  (p. 
95).  If  the  two  gases  are  mixed,  the  molecules  of  hydrogen 
must,  through  frequent  collisions  with  molecules  of  oxygen,  soon 
have  the  same  average  energy  as  the  latter,  and  to  do  this  must 
have,  on  the  average,  four  times  the  velocity  of  the  molecules 
of  oxygen.  What  is  true  of  the  mixed  gases  must  be  true  also 
of  the  gases  when  separate.  Accordingly,  if  we  have  a  porous 
wall  with  fine  openings  separating  the  two  gases,  as  the  hydro- 


DISSOCIATION 


59 


gen  molecules  have  four  times  the  velocity  of  the  oxygen  mole- 
cules and  there  are  the  same  number  in  equal  volumes,  four 
hydrogen  molecules  will  hit  the  openings  and  pass  into  them 
while  one  oxygen  molecule  does  so.  This  ratio  of  four  to  one 
is  title  same  as  the  ratio  of  the  square  roots  of  the  densities, 
Vl6 :  Vl  =  4:1,  the  law  for  diffusion  through  a  porous  wall 
given  above. 

Chemical  Properties  of  Hydrogen.  If  hydrogen  is  brought  to 
the  air,  through  a  glass  tube  drawn  to  a  narrow  opening,  and 
lighted,  it  will  burn  for  an  instant  with  a  pale  blue,  almost  in- 
visible flame,  but  the  color  quickly  changes  to  yellow  from  par- 
ticles of  sodium  or  its  compounds,  which  are  volatilized  from  the 
glass  by  the  heat.  From  a  platinum  jet  the  pure  gas  burns  with 
a  flame  almost  or  quite  invisible  in  daylight.  The  product 
formed  is  water,  as  may  be  shown,  roughly,  by  holding  a  cold 
glass  over  the  flame,  or,  more  accurately,  by  burning  the  gas  for 
some  time,  condensing  the  water  formed  and  determining  its 
freezing  point  and  boiling  point.  The  hydrogen  and  air  used 
should,  of  course,  be  carefully  dried. 

Mixtures  of  air  and  hydrogen  explode  when  ignited,  as  the 
flame  travels  through  such  a  mixture  with  a  very  high  velocity, 
and  both  the  steam  formed  and  the  nitrogen 
of  the  air  are  heated  to  a  high  temperature 
and  expanded  greatly  by  the  heat  of  com- 
bustion. Mixtures  of  oxygen  and  hydrogen 
explode  still  more  violently.  Hydrogen  does 
not  support  the  combustion  of  substances 
which  burn  in  oxygen  (Fig.  21).  It  will 
support  the  combustion  of  oxygen  or  chlo- 
rine. 

Dissociation.  .If  steam  is  passed  through 
a  tube  of  porous  porcelain  A,  Fig.  22,  which 
is  inclosed  in  a  larger  tube  B  of  glazed  porce- 
lain through  which  is  passed  a  current  of 
carbon  dioxide,  introduced  through  C,  while  the  whole  is 
heated  to  a  very  high  temperature  (2000°,  perhaps)  the  steam 


60 


A  TEXTBOOK  OF  CHEMISTRY 


will  be  partly  decomposed  into  oxygen  and  hydrogen,  and, 
since  the  hydrogen  diffuses  through  the  porous  porcelain  four 
times  as  fast  as  the  oxygen,  more  oxygen  than  enough  to  com- 
bine with  the  hydrogen  which  remains  will  stay  in  the  inner 
tube  while  more  hydrogen  will  pass  through  than  enough  to 
combine  with  the  oxygen  which  passes  through.  If  the  cooled 


Fig.  22 

gases  which  are  delivered  at  the  ends  of  the  tube  are  mixed 
and  the  carbon  dioxide  is  absorbed  by  a  solution  of  sodium  hy- 
droxide, it  will  be  found  that  the  gas  which  remains  consists 
of  a  mixture  of  two  volumes  of  hydrogen  with  one  volume  of 
oxygen.  In  this  way  Deville  showed  that  water  can  be  de- 
composed into  oxygen  and  hydrogen  by  heat  alone,  and  that 
the  reaction  between  hydrogen  and  oxygen  is  reversible  : 


A  decomposition  of  this  kind,  when  the  products  of  decomposi- 
tion recombine  on  cooling  or  on  a  reversal  of  the  process  which 
caused  the  decomposition,  is  called  a  dissociation.  Probably  all 
compounds  would  be  decomposed  into  their  elements  at  a  suffi- 
ciently high  temperature,  but  it  is  only  in  those  cases  where  the 
elements  recombine,  on  cooling,  to  form  the  same  compound, 
that  the  decomposition  is  called  a  dissociation. 

*  By  other  methods  the  per  cent  of  dissociation  of  water  has 
been  determined  up  to  2300°,  absolute.  From  the  results  the 
dissociation  at  still  higher  temperatures  may  be  calculated 


OXYHYDROGEN  BLOWPIPE 


61 


approximately.1    The   results   are   as  follows,  at  atmospheric 
pressure : 


ABSOLUTE  TEMPERATURE 

TEMPERATURE  CENTI- 
GRADE 

PER  CENT  OP 
DISSOCIATION 

T 

r° 

1000° 

727° 

0.00003 

1500° 

1227° 

0.022 

2000° 

1727° 

0.59 

2500° 

2227° 

3.98 

2773° 

2500° 

8.12 

3273° 

3000° 

20.0 

4000° 

3727° 

40.5 

The  Oxyhydrogen  Blowpipe.  If  oxygen  and  hydrogen  are 
brought  together  in  such  a  manner  that  they  burn  as  they  come 
in  contact,  an  almost  colorless  flame  having  a  very  high  tempera- 


Fig.  23 

ture  is  produced.     If  steam  could  be  heated  to  very  high  tem- 
peratures by  the  expenditure  of  the  same  quantity  of  energy,  pro- 


1  By  the  formula, 
2  P  (100 


log 


-  i.oo.iu  -^r  -  1000)  -  0.685. 10-7 (T2  -  10002),  in  which  P  is 
the  pressure  in  atmospheres,  T,  the  absolute  temperature,  and  x, 
the  fraction  dissociated.  (Nernst,  Theoretische  Chemie,  6te  Aufl., 
p.  681.) 


62  A  TEXTBOOK  OF  CHEMISTRY 

portionally,  as  that  required  to  heat  it  to  1000°,  the  heat  of  com- 
bustion of  oxygen  and  hydrogen  (p.  65)  is  great  enough  to  heat 
the  steam  produced  to  10,000°,  at  least.  A  little  consideration 
of  the  dissociation  of  water  at  high  temperatures  shows  us  that 
such  an  extreme  temperature  cannot  be  reached,  for  it  is  evident 
that  oxygen  and  hydrogen  cannot  combine  to  produce  heat  at 
a  temperature  at  which  water  largely  decomposes  into  its  ele- 
ments. Since  there  is  a  dissociation  of  40' per  cent  at  3700°,  it  is 
doubtful  if  even  that  temperature  can  be  obtained.  The  tem- 
perature of  an  electric  arc  between  carbon  points  is  estimated 
at  about  3600°.  The  temperature  of  an  open-hearth  steel  fur- 
nace is  only  1500°  to  1700°. 

Iron  wire  will  take  fire  in  the  oxyhydrogen  flame  and  burns 
brilliantly.  Platinum  melts  easily  (1755°)  and  the  flame  has  long 
been  used  in  working  with  this  metal.  A  piece  of  lime  held  in 
the  flame  glows  intensely,  giving  the  light  known  variously  as 
the  oxyhydrogen,  lime  or  Drummond  light.  It  will  be  noticed 
that  any  solid  substance,  which  does  not  volatilize  too  easily, 
gives  an  intense  light  in  the  flame,  while  the  oxyhydrogen  flame 
alone  is  almost  nonluminous. 

Explosions.  Catalysis.  A  mixture  of  oxygen  and  hydrogen 
may  remain  in  a  glass  tube  indefinitely  without  combining  to  an 
extent  that  can  be  measured.  It  is  not  till  a  temperature  of 
300°  is  reached  that  the  gases  combine  rapidly  enough  so  that 
the  rate  of  combination  can  be  measured  after  some  days  or 
weeks.  At  a  temperature  a  little  above  500°,  the  combination 
is  fast  enough  so  that  the  heat  of  combination  raises  the  mix- 
ture to  a  higher  temperature,  which  still  further  accelerates  the 
combination,  and  an  explosion  results.  It  is  characteristic  of 
most  explosions  caused  by  chemical  action  that  the  reaction 
causing  the  explosion  is  accelerated  enormously  by  the  heat 
of  the  reaction.  If  the  mixture  of  oxygen  and  hydrogen  is 
brought  into  contact  with  platinum  in  the  spongy  form  or  with 
platinized  asbestos  1  the  reaction  is  hastened  and  the  gases  will 

1  Prepared  by  moistening  asbestos  with  a  strong  solution  of 
chloroplatinic  acid  and  heating  it. 


OXIDATION.    VALENCE  63 

usually  take  fire  and  burn.  This  is  another  illustration  of  ca- 
talysis and  recalls  the  effect  of  other  metals  on  the  solution  of 
zinc  in  acids.  It  may  be  that  the  two  phenomena  are  closely 
related. 

Oxidation.  Reduction.  If  a  piece  of  copper  is  held  over  a 
flame  so  that  it  is  heated  quite  hot  while  exposed  to  the  air,  it 
will  be  oxidized,  the  surface  changing  to  black  copper  oxide. 
On  holding  the  hot,  oxidized  copper  in  an  atmosphere  of  hydrogen, 
the  black  oxide  will  be  quickly  changed  back  to  metallic  copper. 
This  process  is  called  reduction,  and  hydrogen  is  called  a  reducing 
agent.  It  will  be  seen  from  the  above  that  oxidation  and  reduc- 
tion are  opposite  processes.  The  two  words  are  often  used  in 
chemistry  in  a  much  more  general  sense.  The  addition  of  other 
elements  than  oxygen  is  frequently  called  oxidation,  and  the  re- 
moval of  other  elements,  or  the  substitution  of  hydrogen  for 
other  elements,  or  even  the  addition  of  hydrogen  to  a  com- 
pound, may  be  called  a  reduction. 

Valence.  It  may  have  been  noticed  that  when  sodium  and 
zinc  act  on  water  or  on  hydrochloric  acid,  one  atom  of  the  sodium 
replaces  one  atom  of  hydrogen,  while  one  atom  of  zinc  replaces 
two  atoms.  This  characteristic  of  metals  may  be  made  more 
striking  by  selecting  three  metals  whose  atomic  weights  are  close 
together.  If  23  milligrams  of  sodium,  24  milligrams  of  magne- 
sium and  27  milligrams  of  aluminium  are  allowed  to  act  on  hy- 
drochloric acid  in  such  a  way  that  the  hydrogen  generated  is 
collected  in  separate  tubes,1  it  will  be  found  that  the  sodium  will 
give  about  11  cc.  of  hydrogen,2  the  magnesium  22  cc.  and  the 
aluminium  33  cc.  The  property  of  the  metals  illustrated  here 
is  called  valence.  A  metal  which  replaces  one  atom  of  hydrogen 

1  The  sodium  may  be  placed  in  a  short  lead  tube,  3  mm.  in  diam- 
eter and  closed  at  one  end.     The  mouth  should  be  closed  with  a 
little  cotton  till  ready  for  use.     The  magnesium  and  aluminium  may 
be  weighed  in  small  watch  glasses  about  the  size  of  the  mouths  of 
the  test  tubes  to  be  placed  over  them.     The  aluminium  must  be 
etched  with  a  solution  of  sodium  or  potassium  hydroxide  and  after- 
wards washed  and  dried  before  use. 

2  One  cc.  of    hydrogen  weighs    0.09  mg.     0.09  X  11.1  =  1    nig. 
of  hydrogen  from  23  mg.  of  sodium. 


64 


A  TEXTBOOK  OF  CHEMISTRY 


for  one  atom  of  the  metal  is  called  univalent;  one  which  re- 
places two  atoms  is  called  bivalent;  three,  trivalent;  four, 
quadrivalent.  On  a  somewhat  different  basis,  which  will  be 
discussed  later  (p.  156),  when  other  meanings  of  valence 
are  considered,  some  elements  may  be  quinquivalent,  sexivalent, 
septivalent  or  even  octovalent.  The  valence  of  the  metals  is 


Fig.   24 

such  an  important  characteristic  and  a  knowledge  of  it  is  so 
useful  in  the  writing  of  formulas  that  it  seems  best  to  present 
it  from  the  standpoint  of  replacement  here. 

Some  metals  have  two  or  more  kinds  of  valence  in  different 
compounds.  Thus,  in  the  sense  in  which  valence  is  used  here, 
iron  is  bivalent  in  ferrous  chloride,  FeCl2,  and  ferrous  sulfate, 
FeSO4,  but  trivalent  in  ferric  chloride,  FeCla,  and  ferric  sulfate, 
Fe2(SO4)s.  This  last  formula  illustrates  the  value  of  a  knowledge 
of  valence  in  writing  formulas.  (What  is  the  formula  of  alu- 
minium sulfate?  What  are  the  formulas  of  stannous  sulfate 
and  of  stannic  sulfate  if  tin  is  bivalent  in  stannous  compounds 
and  quadrivalent  in  stannic  compounds  ?) 

From  the  point  of  view  of  the  atomic  theory  the  facts  by 
means  of  which  the  valences  of  elements  are  determined  point 
very  strongly  to  differences  in  the  powers  of  atoms  to  combine 
with  other  atoms.  Thus  in  the  compounds  NaCl,  MgCl2  and 


HYDROGEN  65 

"  Aids  it  seems  that  an  atom  of  sodium  can  hold  one  atom  of  chlo- 
rine in  combination,  an  atom  of  magnesium  can  hold  two  and  an 
atom  of  aluminium,  three.  This  property  of  valence  is  often 
expressed  by  means  of  such  formulas  as  the  following : 

Cl       /Cl 

Na— Cl,  Mg<     ,  Alf-Cl,     H— O— H 
XC1      XC1 

The  lines  are  intended  to  represent  lines  of  force  holding  the 
atoms  in  combination,  and  the  number  of  lines  proceeding  from 
the  symbol  of  an  element  indicates  its  valence. 

We  must  distinguish  sharply  between  the  intensity  of  the  force 
holding  two  atoms  together  and  the  valence  of  the  atoms.  Thus 
the  valence  of  sodium  in  Na — Cl  is  the  same  as  that  of  hydrogen 
in  H — Cl,  but  the  force  which  holds  the  sodium  and  chlorine 
together  is  much  greater  than  that  which  holds  hydrogen  and 
chlorine  together. 

Heat  of  Combustion  of  Hydrogen.  The  heat  generated  when 
2.015  grams  of  hydrogen  combine  with  16  grams  of  oxygen  and 
the  water  formed  is  condensed  to  the  liquid  state  at  18°  is 
68,414  small  calories,  or  34,179  calories  for  one  gram  of  hydro- 
gen (Thomsen).  This  value  is  often  used  in  calculating  the 
heat  of  combustion  of  coal. 

If  2.015  grams  of  hydrogen  combine  with  16  grams  of  oxygen 
at  100°  and  the  water  formed  remains  as  steam,  the  heat  of 
combination  is  only  58,000  calories,  at  constant  pressure,  or 
28,970  calories  for  one  gram.  This  is,  of  course,  the  maximum 
amount  of  heat  which  can  be  obtained  by  burning  hydrogen  un- 
der practical  conditions. 


CHAPTER  V 
WATER,   HYDROGEN   PEROXIDE 

Analysis.  Synthesis.  The  two  methods  by  which  the  com- 
position of  a  substance  is  determined  are  by  analysis,  the  sepa- 
ration of  the  substance  into  the  elements  of  which  it  is  com- 
posed, and  by  synthesis,  the  putting  together  of  the  elements  to 
form  the  compound.  In  analysis  it  is  comparatively  seldom  that 
the  elements  are  separated  in  the  free  state.  Thus,  in  order  to 
determine  the  amount  of  hydrogen  in  an  organic  compound, 
such  as  sugar,  the  substance  is  burned  and  the  water  formed  is 
collected  and  weighed.  Knowing  what  per  cent  of  hydrogen 
is  contained  in  water,  it  is  easy  to  calculate  the  amount  of  hydro- 
gen contained  in  the  compound.  Either  an  analysis  or  a  synthe- 
sis may  be  qualitative,  giving  simply  the  elements  which  are 
present,  or  quantitative,  giving  the  quantity  or  per  cent  of  each 
element. 

Qualitative  Analysis  and  Synthesis  of  Water.  The  prepara- 
tion of  hydrogen  by  passing  steam  over  red-hot  iron  is  a  qualita- 
tive analysis  of  water.  In  order  to  make  the  analysis  complete 
it  would  be  necessary  to  show  that  hydrogen  and  the  magnetic 
oxide  of  iron  are  the  only  products  of  the  action,  that  the  density 
and  properties  of  the  hydrogen  are  the  same  as  those  of  hydrogen 
prepared  in  other  ways  and  that  the  density  and  properties  of 
the  magnetic  oxide  formed  are  the  same  as  those  of  the  oxide 
formed  by  burning  iron  in  oxygen. 

The  experiment  showing  that  water  is  formed  when  dry  hy- 
drogen burns  in  air  or  in  oxygen  is  a  qualitative  synthesis  of 
water. 

Quantitative  Synthesis  of  Water  by  Volume.  The  determina- 
tion of  the  composition  of  water  by  volume  may  be  made  in  an 
instrument  called  a  eudiometer,  a  tube  graduated  usually  to 

66 


COMPOSITION  OF  WATER 


67 


tenths  of  a  cubic  centimeter  by  means  of  fine  lines  etched  on 
the  surface.  For  the  synthesis  of  water  two  platinum  wires 
must  be  sealed  in,  near  the  closed  end.  A  capacity  of  50  cc.  is 
suitable  for  the  experiment  to  be  described.  Such  a  eudiometer 
is  carefully  dried  and  filled  with  mercury  and  10-12  cc.  of  dry 
oxygen  introduced.  The  volume  of  the  gas  is  then  accurately 
measured  and  the  height  of  the  mercury  in  the  eudi- 
ometer above  the  mercury  in  the  reservoir,  the  tem- 
perature and  the  reading  of  the  barometer  are  noted. 
From  these  measurements  the  volume  of  the  oxygen 
under  standard  conditions  is  calculated.  Enough 
dry  hydrogen  to  give  a  total  volume  of  25-28  cc.  is 
then  introduced  and  these  measurements  repeated. 
The  hydrogen  should  be  in  excess,  but,  as  the  pres- 
sure is  greater,  the  total  volume 
need  not  be  three  times  the  vol- 
ume of  the  oxygen.  The  eudiom- 
eter is  then  clamped  firmly,  with 
a  piece  of  sheet  rubber  placed 
under  its  mouth,  and  the  mix- 
ture exploded  by  passing  an  elec- 
tric spark  between  the  platinum 
points  by  means  of  an  induction 
coil.  After  cooling,  the  volume 
of  hydrogen  remaining  is  meas- 
ured as  before,  except  that  the 
water  formed  by  the  explosion 
remains  partly  as  vapor  in  the 


Steam 


=18 


=20 


.-22 


=24 


-23 


=32 


=36 


Fig.  25 


hydrogen  and  the  pressure  of  the  hydrogen  is 
to  be  found  by  subtracting  from  the  reading 
of  the  barometer  the  height  of  the  mercury 
in  the  eudiometer  plus  the  pressure  of  vapor 
of  water  for  the  temperature  which  is  read 
(see  p.  75).     After  correction  of   the  three  volumes   of  gas  to 
standard  conditions  the  proportion  by  volume  in  which  the  gases 
have  united  may  be  calculated. 


68  A  TEXTBOOK  OF  CHEMISTRY 

By  placing  a  tube  over  the  eudiometer  (Fig.  26)  and  passing 
steam  through  it,  the  water  formed  in  the  experiment  may  be 
converted  into  steam  and  the  volume  of  the  excess  of  hydrogen 
plus  the  steam  determined  and  from  this  the  volume  of  the  steam 
calculated. 

*  The  results  of  Morley's  exceedingly  careful  experiments 
(Amer.  J.  Sci.  41,  220  and  276 ;  Chem.  News,  63,  218)  show  that 
when  oxygen  and  hydrogen  are  measured  in  tubes  the  ratio  of 
the  volumes  which  combine  is,  O :  H  =  1 :  2.0002.  Curiously 
enough  Scott  (Phil.  Trans.  184,  A,  543  (1893))  has  found  that 
when  the  gases  are  measured  in  globes  the  ratio  is  O :  H  = 
1 :  2.0025.  It  seems,  therefore,  that  the  same  quantity  of  gas 
may  fill  a  different  volume  when  measured  in  a  tube  from  what 
it  does  when  measured  in  a  globe,  but  no  one  has  proved  this  by 
direct  experiment. 

The  volume  of  the  steam  is  very  nearly  the  same  as  the  volume 
of  the  hydrogen  which  goes  to  form  it.  We  can  express  the  re- 
lation by  the  following  diagram  : 


II 


1  vol.  oxygen  2  vols.  hydrogen  2  vols.  steam 

Composition  of  Water  by  Weight.  From  the  composition  of 
water  by  volume  and  the  weights  of  one  liter  of  each  gas  we  may 
calculate  the  composition  by  weight.  This  gives : 

O  :  H  =  1.429  :  2.0025  X  0.08987 
or  16:2  X  1.0075 

The  Unit  for  Atomic  Weights.  The  quantity  of  hydrogen 
combining  with  16  parts  of  oxygen  is  given  because  an  atomic 
weight  of  16  has  been  assigned  to  oxygen,  somewhat  arbitrarily, 
as  a  basis  for  comparison  with  all  other  atomic  weights.  Hy- 
drogen with  an  atomic  weight  of  one  was  originally  chosen  as 
the  unit  for  atomic  weights.  For  70  or  80  years  it  was  sup- 
posed, on  the  basis  of  inaccurate  determinations  of  the  composi- 
tion of  water,  that  the  atomic  weight  of  oxygen  was,  on  that 


COMPOSITION  OF  WATER  69 

basis,  almost  exactly  16  (or  8).  When  the  composition  of  water 
was  finally  determined  more  accurately,  it  was  decided  by  the 
majority  of  chemists  that  it  is  better  to  make  oxygen,  with  an 
atomic  weight  of  16,  the  basis  for  all  other  atomic  weights,  rather 
than  to  make  such  large  changes  as  would  be  necessary  in  many 
of  the  common  atomic  weights,  if  hydrogen  were  retained  as 
the  unit.  A  few  chemists,  however,  still  prefer  hydrogen  as  the 
unit. 

Determination  of  the  Composition  of  Water  by  the  Use  of 
Copper  Oxide.  The  first  moderately  accurate  determination  of 
the  composition  of  water  was  made  by  the  Swedish  chemist, 
Berzelius,  in  1819.  He  weighed  a  quantity  of  copper  oxide  in 
a  glass  bulb,  heated  it,  passed  dry  hydrogen  through  the  bulb, 
and  collected  and  weighed  the  water  formed.  The  copper  oxide 
was  reduced  to  metallic  copper,  and  the  loss  of  weight  gave  the 
weight  of  oxygen  which  had  been  converted  into  water.  The 
difference  between  the  weight  of  the  water  collected  and  the 
weight  of  oxygen  taken  from  the  copper  oxide  gave  the  weight 
of  the  hydrogen.  A  number  of  years  later  (1842)  a  French 
chemist,  Dumas,  carried  out  an  elaborate  series  of  experiments 
by  the  same  method,  with  the  apparatus  shown  in  Fig.  27. 
The  hydrogen  was  generated  in  the  large  bottle  and  passed 
through  a  series  of  tubes  to  purify  and  dry  it.  It  was  then 
passed  through  the  bulb  containing  the  copper  oxide  and  the 
water  formed  was  collected,  partly  in  a  bulb,  and  partly  in  dry- 
ing tubes.  The  atomic  weight  of  hydrogen  calculated  from  the 
results  of  19  experiments  with  this  apparatus  is  1.0025.1  For 
some  unknown  reason  the  result  is  too  low  by  about  one  part 
in  200.  About  50  years  later  the  copper  oxide  method  was 
modified  by  the  author,  who  used  the  apparatus  shown  in  Fig. 
28.  After  placing  some  copper  oxide  in  the  bulb  A  and  ex- 

1  The  student  is  not,  of  course,  expected  to  remember  these 
various  values.  A  very  brief  description  of  4  out  of  some  16  deter- 
minations of  the  composition  of  water  is  given  as  an  illustration  of  the 
amount  of  labor  which  has  been  expended  on  the  determination  oi 
atomic  weights  and  also  to  illustrate  how  successive  workers  attain, 
sometimes,  a  closer  approximation  to  the  truth. 


70 


A  TEXTBOOK  OP  CHEMISTRY 


COMPOSITION  OF  WATER  71 

hausting  it  with  a  good  mercury  air  pump  the  apparatus  was 
weighed.  It  was  then  connected  at  C  with  an  apparatus  fur- 
nishing pure  hydrogen,  the  bulb  A  was  heated  in  an  air  bath  and 
the  tube  B  was  cooled.  The  stopcock  E  permitted  the  hydrogen 
to  pass  out  at  first  through  D  so  that  no  air  should  enter  the  bulb 
from  the  connecting  tubes.  On  admitting  hydrogen  to  the  bulb 
it  was  converted  into  water  by  the  copper  oxide  and  the  water 
was  condensed  in  B.  After  from  one  to  two  grams  of  hydro- 


Fig.  28 

gen  had  been  converted  into  water  in  this  way  the  stopcock  was 
closed  and  the  apparatus  cooled  and  weighed.  The  gain  in 
weight  was  the  weight  of  the  hydrogen  which  had  entered.  The 
apparatus  was  then  connected  with  a  tube  into  which  the  water 
formed  could  be  driven  by  warming  B  and  A.  The  loss  in  weight 
of  the  apparatus  gave  the  weight  of  the  oxygen  which  had  been 
taken  from  the  copper  oxide.  The  water  was  also  collected  and 
weighed.  Twenty-four  determinations,  partly  by  this  method, 
partly  by  another  which  need  not  be  described  here,  gave 
1.00787  as  the  atomic  weight  of  hydrogen. 

Determination  of  the  Composition  of  Water  by  weighing 
Oxygen  and  Hydrogen.  For  more  than  twelve  years  Professor 
Morley  worked  at  Cleveland  on  the  composition  of  water  by 
volume,  on  the  determination  of  the  weights  of  oxygen  and  hy- 
drogen gases  and  finally  on  the  composition  of  water  by  weight. 
The  weights  of  one  liter  of  hydrogen  and  of  oxygen  which  have 
been  given  are  the  values  determined  in  this  long,  classical  in- 


72 


A  TEXTBOOK  OF  CHEMISTRY 


vestigation.  The  composition  of  water  by  weight  was  also  de- 
termined in  the  apparatus  shown  in  Fig.  29.  Into  this  apparatus 
were  brought  hydrogen  from  a  tube  containing  metallic  palla- 
dium, in  which  it  had  been  absorbed,  and  oxygen  from  globes  in 
which  it  had  been  weighed.  The  water  formed  was  frozen  in 
the  bottom  of  the  apparatus  and  was  weighed 
at  the  end  of  the  experiment.  From  twelve 
experiments  the  atomic  weight  of  hydrogen  is 
calculated  as  1.00762.  The  atomic  weight  of 
hydrogen  which  is  now  used  (J.  Am.  Chem. 
Soc.  31,  1)  is  1.0078.  This  is  probably  not  in 
error  by  so  much  as  one  part  in  5000.  For 
ordinary  calculations  the  value  is  rounded  off 
to  1.008,  or,  frequently,  to  1.01. 

Properties  of  Water.  Pure  water  appears, 
ordinarily,  to  be  colorless  and  transparent,  but 
light  transmitted  through  a  layer  of  water  some 
meters  in  thickness  has  a  blue  color.  Water 
is  the  only  substance  for  which  we  have  three 
names  according  as  it  is  in  the  form  of  a  solid, 
liquid  or  gas.  Water  has  a  maximum  density 
at  4°.  If  either  cooled  or  heated  from  that 
temperature,  it  expands.  For  this  reason  in 
the  fall  and  winter  large  bodies  of  water  cool 
by  convection,  that  is  by  the  sinking  of  the 
cooler,  heavier  water  on  the  surface,  till  a  temperature  of  4°  is 
reached.  On  further  cooling  the  water  grows  lighter  again  and 
the  colder,  lighter  water  floats  on  the  surface,  protecting  the 
warmer  water  beneath  from  further  rapid  cooling.  The  ice  which 
finally  forms  has  a  density  of  only  0.92  (accurately  0.91674)  and 
continues  to  float  on  the  surface.  The  density  of  water  at  dif- 
ferent temperatures  is  given  in  the  following  table  : l 

1  This  table  is  useful  especially  for  determining  the  capacity  of 
burettes  or  flasks  by  weighing  the  water  which  they  contain.  Thus 
it  is  seen  that  the  water  which  will  fill  1  cc.  at  20°  weighs  0.99823 
gram.  This  is,  however,  when  weighed  in  a  vacuum.  If  weighed 
with  brass  weights  in  air,  the  apparent  weight  will  be  0.00105  gram. 


Fig.  29 


PROPERTIES  OF  WATER 


73 


TEMPERATURE 

DENSITY 

TEMPERATURE 

DENSITY 

0 

0.99987 

20 

0.99823 

1 

0.99993 

21 

0.99802 

2 

0.99997 

22 

0.99780 

3 

0.99999 

23 

0.99756 

4 

1.00000 

24 

0.99732' 

5 

0.99999 

6 

0.99997 

25 

0.99707 

7 

0.99993 

30 

0.99567 

8 

0.99988 

35 

0.99406 

9 

0.99981 

40 

0.99224 

45 

0.99024 

10 

0.99973 

50 

0.98807 

11 

0.99963 

55 

0.98573 

12 

0.99952 

60 

0.98324 

13 

0.99940 

65 

0.98059 

14 

0.99927 

70 

0.97781 

75 

0.97489 

15 

0.99913 

80 

0.97183 

16 

0.99897  ' 

85 

0.96865 

17 

0.99880 

90 

0.96534 

18 

0.99862 

95 

0.96182 

19 

0.99843 

100 

0.95838 

less,  or  0.99718  gram.  The  following  table,  which  gives  the  appar- 
ent weight  of  one  liter  of  water  weighed  with  brass  weights  in  air 
and  the  corresponding  correction  to  volume,  is  still  more  conven- 
ient. 

Apparent  weight  of  one  liter  of  water  in  air  and  corrections  to  be 
added  to  the  apparent  weight  of  one  liter  of  water  to  find  the  true 
volume  in  cubic  centimeters. 


TEMPERA- 
TURE 

GRAMS 

CORRECTION 

TEMPERA- 
TURE 

GRAMS 

CORRECTION 

15° 

998.05 

2.07  CC. 

23° 

996.53 

3.40  CC. 

16° 

997.90 

2.20  cc. 

24° 

996.29 

3.61  cc. 

17° 

997.74 

2.34  cc. 

25° 

996.04 

3.83  cc. 

18° 

997.56 

2.49  cc. 

26° 

995.79 

4.06  cc. 

19° 

997.38 

2.65  cc. 

27° 

995.52 

4.31  cc. 

20° 

997.18 

2.82  cc. 

28° 

995.24 

4.56  cc. 

21° 

996.97 

3.00  cc. 

29° 

994.96 

4.82  cc. 

22° 

996.76 

3.19  cc. 

30° 

994.66 

5.08  cc. 

74 


A  TEXTBOOK  OF  CHEMISTRY 


Heat  of  Fusion  and  Vaporization.  If  heat  is  applied  to  a 
kilogram  of  ice  at  the  freezing  point,  0°,  it  will  absorb  79  large 
calories  in  melting ;  that  is,  if  the  same  amount  of  heat  is  applied 
to  a  kilogram  of  ice  and  a  kilogram  of  water,  both  of  them  at  0°, 
when  the  ice  is  melted  and  still  at  0°  the  other  water  will  be  at 
a  temperature  of  79°.  If  a  kilogram  of  water  at  the  boiling 
point,  100°,  is  heated  till  it  is  all  converted  into  steam  at  the 
same  temperature  and  at  atmospheric  pressure  it  will  absorb 
536  large  calories,  or  enough  to  warm  5.36  kilograms  of  water 
from  the  freezing  point  to  the  boiling  point.  The  steam  formed 
will  fill  a  space  of  about  1700  liters,  while  the  liquid  water  fills 
only  one  liter. 

The  heat  which  is  absorbed  in  the  melting  of  ice  and  vaporiza- 
tion of  water  was  formerly  called  latent  heat  because  it  seems  to 
disappear,  but  this  expression  is  not  as  fre- 
quently used  now  as  it  was  some  years  ago. 
The  heat  of  vaporization  varies  with  the  tem- 
perature, being  greater  at  lower  and  less  at 
higher  temperatures. 

Vapor  Pressure  of  Water.  If  two  dry 
tubes,  about  800  mm.  long,  are  filled  with 
mercury  and  inverted,  the  mercury  will  fall 
until  the  weight  of  the  mercury  in  the  tubes 
is  the  same  as  the  weight  of  a  column  of  air 
of  the  same  cross  section  and  reaching  to  the 
top  of  the  atmosphere.  In  other  words,  the 
mercury  will  stand  at  the  same  height  as  the 
mercury  in  a  barometer.  If  a  drop  of  water 
is  introduced  into  one  of  the  tubes,  the  mer- 
cury in  the  tube  will  fall  and  remain  at  a 
position  lower  than  that  in  the  dry  tube.  If 
the  tube  containing  the  water  is  warmed,  the 
mercury  will  fall  further,  if  cooled,  it  will  rise  higher  and  for  each 
temperature  there  will  be  a  definite  difference  between  the 
heights  of  the  mercury  in  the  two  tubes.  It  is  evident  that  this 
must  be  caused  by  the  fact  that  a  part  of  the  water  in  the  tube 


Fig.  30 


VAPOR  PRESSURE   OF  WATER 


75 


is  converted  into  a  vapor  or  gas  and  exerts  a  pressure  on  the 
mercury,  partially  balancing  the  pressure  of  the  air.  This  pres- 
sure is  called  the  pressure  of  water  vapor,  or  sometimes,  and  less 
correctly,  the  aqueous  tension.  It  is  given  for  different  tempera- 
tures in  the  following  table : 

VAPOR  PRESSURE  OF  ICE   AND  WATER 


TEMPERA- 
TUBES 

PRESSURE  IN 
MILLIMETERS 
OF  MERCURY 

TEMPERA- 
TURES 

PRESSURE 

IN  MM. 

TEMPERA- 
TURES 

PRESSURE 

IN   MM. 

-10° 

2.0 

27° 

26.5 

100.5° 

773.7 

-  5° 

3.0 

28° 

28.1 

101.0° 

787.6 

2° 

3.9 

29° 

29.8 

-   1° 

4.2 

30° 

31.6 

±  0° 

4.6 

31° 

33.4 

1° 

4.9 

32° 

35.4 

2° 

5.3 

33° 

37.4 

3° 

5.7 

34° 

39.6 

PRESSURE  IN 

4° 

6.1 

35° 

Af\O 

41.9 

er  er  A 

ATMOSPHERES 

5° 
6° 

6.5 

7.0 

40 
50° 

55.0 
92.2 

111.7° 

1.5 

7° 

7.5 

60° 

149.2 

120.6° 

2 

8° 

8.0 

70° 

233.8 

127.8° 

2.5 

9° 

8.6 

80° 

355.5 

133.9° 

3 

10° 

9.2 

90° 

526.0 

144.0° 

4 

11° 

9.8 

95° 

634.0 

159.2° 

6 

12° 

10.5 

96° 

657.7 

170.8° 

8 

13° 

11.2 

97° 

682.1 

180.3° 

10 

14° 

11.9 

98°  • 

707.3 

188.4° 

12 

15° 

12.7 

99° 

733.2 

195.5° 

14 

16° 

13.6 

99.1° 

735.9 

201.9° 

16 

17° 

14.5 

99.2° 

738.5 

207.7° 

18 

18° 

15.4 

99.3° 

741.2 

213.0° 

20 

19° 

16.4 

99.4° 

743.9 

'224.7° 

25 

20° 

17.4 

99.5° 

746.5 

"    21° 

18.5 

99.6° 

749.2 

22° 

19.7 

99.7° 

751.9 

23° 

20.9 

99.8° 

754.6 

24° 

22.2 

99.9° 

757.3 

25° 

23.5 

100.0° 

760.0 

26° 

25.0 

100.1° 

762.7 

100.2° 

765.5 

100.3° 

768.2 

100.4° 

770.9 

76  A  TEXTBOOK  OF  CHEMISTRY 

Equilibrium.  So  long  as  water  remains  in  the  liquid  form  in  the 
barometer  tube  described  in  the  last  paragraph,  the  volume  of 
the  tube  above  will  have  no  effect  on  the  vapor  pressure  of  the 
water.  This  relation,  which  is  very  important,  will  be  clearer 
from  Fig.  31.  Suppose  that  the  cylinder  contains  water  in  the 
bottom  with  vapor  of  water  above  it  and  that  it  is  fitted  with  an 
air-tight  piston.  If  the  piston  is  raised,  the  vapor,  which  acts 
like  any  other  gas  in  this  respect,  will  immediately  expand  and 
fill  this  additional  space,  and  the  pressure  will  be  momentarily 
lowered.  Immediately,  however,  some  of  the  water 
will  evaporate,  and  this  evaporation  will  continue  till 
the  pressure  is  the  same  as  before,  provided  that  the 
water  is  kept  at  the  same  temperature.  If  the  piston 
is  pressed  down,  the  reverse  operation  will  occur.  The 
vapor  will  be  momentarily  compressed  and  the  pressure 
increased,  but  some  of  the  vapor  will  immediately  con- 
dense to  water  and  in  this  way  the  pressure  will  fall  to 
its  original  value.  When  two  or  more  forms  of  a  sub- 
stance are  related  in  this  way,  they  are  said  to  be  in 
equilibrium,  that  is,  they  are  so  balanced  against  each 
other  that  any  change  in  temperature  or  pressure  will 
cause  a  partial  conversion  of  one  form  into  the  other. 
It  is  not  necessary  to  suppose  that  the  same  molecules 
Fig.  31  of  water  are  always  in  the  form  of  vapor  in  such  a  case. 
On  the  contrary,  it  seems  more  probable  that  some  of 
the  water  particles  are  all  of  the  time  leaving  the  liquid  and  pass- 
ing into  the  vapor  and  that  molecules  of  vapor  are  constantly 
leaving  the  vapor  and  passing  back  into  the  liquid.  When  the 
two  are  in  equilibrium,  just  as  many  molecules  must  pass  in  ohe 
direction  as  in  the  other  within  a  given  time. 

Effect  of  Water  Vapor  on  the  Pressure  of  a  Gas.  If  a  small 
glass  bulb  filled  with  water  (Fig.  32)  is  placed  in  a  bottle  filled 
with  dry  air  and  closed  with  a  stopper  bearing  a  manometer 
to  show  the  pressure  within  the  bottle,  on  breaking  the  bulb 
it  will  be  seen  that  the  pressure  within  the  bottle  increases  and 
after  some  time  the  increase  will  be  almost  exactly  equal  to  the 


PHASES 


77 


vapor  pressure  of  water  at  the  temperature  of  the  experiment. 
It  is  evident  from  this  that  as  the  water  evaporates  the  vapor 
diffuses  into  the  air  above  just  as  any  other  gas  would  do,  and 
as  it  does  so  it  adds  its  pressure  to  that  of  the  air,  in  accordance 
with  the  law  of  partial  pressures,  that  each  gas  in  a  mixture 
exerts  the  same  pressure  as  though 
it  were  present  alone  (Dalton's  law 
of  partial  pressures). 

The  experiment  suggests  the 
proper  method  of  finding  the  vol- 
ume which  a  quantity  of  dry  gas 
would  fill  under  standard  condi- 
tions, when  the  gas  has  been  meas- 
ured in  a  moist  condition.  The 
actual  pressure  exerted  by  the  gas 
is  less  than  the  apparent  pressure 
by  the  pressure  exerted  by  the 
vapor  of  water  at  the  given  tem- 
perature. For  this  reason  in  the 
experiment  described  on  p.  67 
the  direction  was  given  to  subtract 
from  the  reading  of  the  barometer 
both  the  height  of  the  mercury 
in  the  eudiometer  and  the  pressure  of  water  vapor,  when 
the  corrected  volume  of  the  moist  hydrogen  was  to  be 
found. 

Phases.  Degrees  of  Freedom.  Water  may  exist  in  the  three 
forms  of  ice,  water  and  vapor.  These  three"  forms  of  the  same 
substance  are  called  phases.  As  long  as  only  one  phase  is  present 
we  may  change  either  the  temperature  or  pressure  or  both  at 
will,  and  in  order  to  know  the  condition  of  the  phase  we  must 
know  both  the  temperature  and  the  pressure.  The  system  is 
said  to  have  two  degrees  of  freedom  and  the  system  is  called 
divariant. 

When  two  phases  are  present,  any  change  in  the  temperature 
will  cause  a  corresponding  change  in  the  pressure,  and  as  long  as 


Fig.  32 


78  A  TEXTBOOK  OF  CHEMISTRY 

the  temperature  is  fixed  the  pressure  cannot  be  changed  without 
the  disappearance  of  one  of  the  phases.  Or,  as  long  as  the  pres- 
sure is  fixed,  the  temperature  cannot  be  changed.  Thus  for 
water  and  vapor  at  a  given  temperature  an  increase  of  the  volume 
does  not  decrease  the  pressure,  but,  instead  of  this,  causes  some 
of  the  water  to  change  to  vapor ;  and  it  is  only  when  the  liquid 
phase  disappears  that  a  further  increase  in  the  volume  causes 
a  decrease  in  the  pressure.  In  the  same  way,  if  water  and  ice 
are  present,  an  increase  in  the  pressure  will  cause  some  of  the 
ice  to  melt  and  the-  temperature  will  fall,  and  for  every  pressure 
there  will  be  a  corresponding  temperature  at  which  ice  and 
water  can  exist  together.  To  know  the  condition  of  such  a 
system  of  two  phases  we  need  give  only  one  factor.  If  we  know 
the  temperature,  the  pressure  is  fixed ;  or  if  we  know  the  pressure, 
the  temperature  is  fixed  by  the  properties  of  the  substance. 
Such  a  system  has  only  one  degree  of  freedom  and  is  called  uni- 
variant. 

When  the  three  phases,  water,  ice  and  vapor,  are  present,  it 
is  impossible  to  change  either  the  temperature  or  the  pressure 
without  the  disappearance  of  one  of  the  phases.  Such  a  system 
has  no  degree  of  freedom  and  is  called  invariant. 

The  relations  between  temperature  and  pressure  for  the  three 
phases  of  water  may  be  seen  clearly  from  the  diagram  (Fig.  33). 
If  water  and  vapor  are  present,  the  relation  between  temperature 
and  pressure  is  fixed  by  the  line  OA.  For  water  and  ice  the  re- 
lation is  fixed  by  the  line  OC,  from  which  it  is  apparent  that  an 
increase  in  the  pressure  lowers  the  melting  point,  though  very 
slowly.  The  relation  for  ice  and  vapor  is  fixed  by  the  line  OB. 
The  line  OB'  gives  the  relation  for  vapor  and  supercooled  water, 
the  vapor  pressure  being  slightly  greater  than  that  of  ice  at  the 
same  temperature.  The  line  OBf  represents  a  condition  of  un- 
stable equilibrium,  and  if  a  little  ice  is  introduced,  the  pressure  will 
fall  or  the  temperature  will  rise  to  the  line  OB,  if  all  of  the  water 
freezes,  or  both  pressure  and  temperature  will  rise  to  the 
point  O,  which  is  the  invariant  point,  called  also  the  triple 
point. 


SOLUTIONS 


79 


*  The  temperature  of  the  system  at  the  triple  point  will  be 
H-  0.0073,  since  the  0°  of  our  scale  is  determined  by  freezing  water 
under  atmospheric  pressure  and  a  pressure  of  one  atmosphere 
lowers  the  freezing  point  0.0073°.  The  vapor  pressure  of  water 
at  the  freezing  point  is  only  4.6  mm. 

Water  as  a  Solvent.  Solutes.  If  salt  or  sugar  is  placed  in 
water,  it  quickly  disappears  and  a  homogeneous  liquid  is  obtained, 
which  we  call  a  solution.  Any  substance  which  passes  into  solu- 
tion in  this  manner  is  called  a  solute.  The  liquid  in  which  the 


|4. 6mm. 


Ice 


Vapor 


-10 


•*  O.O073 
TEMPERATURE 

Fig.  33 


-HO 


solute  dissolves  is  called  a  solvent.  Solutes  may  be  either  solids, 
liquids  or  gases,  and  they  vary  very  greatly  in  their  degrees  of 
solubility.  Thus  one  liter  of  water  will  dissolve  at  20°  670  grams 
of  sugar  or  358  grams  of  common  salt,  but  it  will  dissolve  only 
0.00153  gram  of  silver  chloride.  No  satisfactory  reason  for  such 
differences  has  been  discovered,  though  many  empirical  relations 
between  the  composition  of  substances  and  their  solubility  are 
known. 

Some  substances,  as  alcohol  or  sulfuric  acid,  dissolve  in  water 
in  all  proportions,  but  others  will  dissolve  only  up  to  a  definite 
limit.  When  a  solution  can  remain  in  contact  with  the  solute 


80 


A  TEXTBOOK  OF  CHEMISTRY 


without  taking  up  any  more,  it  is  said  to  be  saturated.  The  solid, 
liquid  or  gaseous  phase  of  the  pure  solute  is  then  in  equilibrium 
with  the  solution  very  much  as  vapor  of  water  is  in  equilibrium 
with  liquid  water. 

The  solubility  of  salts  usually  increases  with  the  temperature, 
but    there    are    some-  exceptions,    and  the    rate    of    increase 
varies  very  greatly,  as  will  be  apparent  from  the  accompany- 
ing   diagram     (Fig. 

240 1 1 1 1 1 T? — rr~i 1 ,    34). 

If  a  warm  satu- 
rated solution  of  a 
salt  which  is  more 
soluble  in  warm 
than  in  cold  water 
allowed  to  cool 


tO 
O 


O 
O 


QO 
O 


Ci 
O 


n 

^ 
O 


£  120 


3  100 


a 


s 

out  of  contact  with 
the  solid  phase,  a 
supersaturated  solu- 
tion may  usually  be 
obtained.  The  in- 
troduction of  a  little 
of  the  solid  will 
start  the  separation 
of  the  solid  phase, 
and  after  a  short 
time  the  solution 
will  assume  the 
normal,  saturated 
condition.  In  a 
similar  manner  still 

water  may  be  cooled  below  its  freezing  point  or  a  vapor  may 
be  cooled  or  compressed  below  the  point  at  which  a  part  would 
ordinarily  exist  in  the  liquid  form.  Such  a  system  is  always 
unstable,  somewhat  after  the  analogy  of  a  pyramid  standing 
on  its  apex,  and  can  only  occur  in  the  absence  of  the  solid  or 
liquid  phase  which  should  normally  be  present. 


20°     40° 


60°     80°    100°   120°  140°  160 
Temperature 

Fig.  34 


SOLUTIONS  81 

Chemical  Activity  in  Solutions.  Metathesis.  If  common  salt, 
NaCl,  and  silver  nitrate,  AgNO3,  are  powdered  and  mixed  to- 
gether, there  will  be  no  apparent  action ;  but  if  each  is  dissolved 
in  water  and  the  solutions  mixed,  there  will  be  formed  immedi- 
ately a  white  precipitate  of  silver  chloride ;  and  if  the  solution 
is  filtered  from  the  precipitate  and  the  filtrate  evaporated, 
crystals  of  sodium  nitrate  may  be  obtained : 

NaCl        +        AgNO3     =     AgCl    +     NaNO3 

Sodium  Chloride  Silver  Silver  Sodium 

(common  salt)  Nitrate  Chloride  Nitrate 

Hundreds  of  illustrations  of  similar  reactions  which  do  not 
occur  readily  between  the  solid  substances  but  which  take  place 
easily  in  solutions  might  be  given.  This  reactivity  of  substances 
in  solution  is  evidently  in  part  because  they  are  brought  into  inti- 
mate contact,  since  no  combination  can  take  place  except  be- 
tween substances  which  are  touching  each  other.  But  this  does 
not  appear  to  be  the  only  reason.  In  very  many  cases  when 
clear,  sharp-cut  reactions  occur,  each  compound  separates,  as 
here,  into  the  metal  and  an  acid  radical.  If  we  subject  these 
same  compounds  in  solution  to  the  influence  of  an  electrical 
current,  the  metal  will  travel  toward  the  cathode  through  the 
solution  while  the  acid  radical  will  travel  toward  the  anode.  It 
seems,  therefore,  that  solution  in  some  way  loosens  the  combina- 
tion between  the  ions  so  that  they  can  very  readily  enter  into 
new  combinations. 

A  reaction  of  the  sort  just  considered  is  called  a  double  decom- 
position or  metathesis. 

Hydrates,  Deliquescence,  Efflorescence.  Many  salts  when 
they  separate  in  crystals  from  solution  do  so  in  combination  with 
a  definite  quantity  of  water.  Thus  crystals  of  copper  sulfate 
have  the  composition  CuSO4.5  H2O ;  crystals  of  sodium  sulfate, 
the  composition  Na2SO4.10H2O.  Such  compounds  usually 
decompose  rather  easily  into  water  and  the  anhydrous  salt,  and 
the  water  is  spoken  of  as  water  of  hydration,  and  the  compounds 


82  A  TEXTBOOK  OF  CHEMISTRY 

are  called  hydrates.1  If  the  hydrate  is  placed  above  the  mercury 
in  a  barometer  tube  it  will  decompose,  giving  off  water  vapor  till  a 
definite  vapor  pressure  is  reached.  This  vapor  pressure  will  vary 
greatly  for  different  hydrates  and  will  increase  with  the  tem- 
perature as  the  vapor  pressure  of  water  does.  Thus  the  vapor 
pressure  of  the  hydrate  CuSO4.5  H2O  is  12.5  mm.  at  30°,  the 
vapor  pressure  of  CaCl2.H2O  is  only  3.1  mm.,  the  vapor  pressure 
of  Na2SO4.10  H2O  is  25.3  mm.  and  the  vapor  pressure  of  pure 
water  is  31.6  mm.  If  an  anhydrous  salt  like  calcium  chloride  is 
exposed  to  air  containing  enough  moisture  so  that  the  pressure 
of  the  water  vapor  in  it  exceeds  3.1  mm.  at  30°,  water  will  be 
absorbed  and  the  hydrate  will  be  formed.  In  this  case  even  a 
concentrated  solution  of  calcium  chloride  has  so  low  a  vapor 
pressure  that  it  is  exceeded  by  that  of  the  moisture  in  ordinary 
air.  Accordingly  calcium  chloride  when  exposed  to  the  air 
absorbs  moisture  and  finally  dissolves  in  the  water  absorbed. 
Such  a  salt  is  said  to  be  deliquescent. 

On  the  other  hand,  if  the  hydrate,  Na2SO4.10  H2O,  is  ex- 
posed to  air  in  which  the  vapor  pressure  of  the  water  which  it 
contains  is  less  than  25.3  mm.  at  30°  the  salt  will  decompose  and 
lose  water  to  the  air.  As  it  does  so  it  will  fall  to  a  fine  powder  or 
flour.  Salts  of  this  type  are  called  efflorescent. 

Natural  Waters.  The  water  which  is  found  in  nature  is  never 
pure,  the  purest  being  rain  water  falling  in  the  open  country 
after  it  has  been  raining  for  some  time,  or  water  obtained  by 
melting  the  ice  from  a  pure,  fresh-water  lake.  Even  such  water 

1  The  term  water  of  crystallization,  which  is  still  used  by  many 
authors,  is  not  as  appropriate.  During  the  first  half  of  the  nine- 
teenth century  sodium  hydroxide,  NaOH,  and  calcium  hydroxide, 
Ca(OH)2,  were  called  sodium  hydrate  and  calcium  hydrate,  and  their 
formulas  were  written  in  a  form  which  with  modern  atomic  weights 
would  be  Na2O.H2O  and  CaO.H?O.  This  older  use  of  the  word 
hydrate,  which  still  clings  to  the  literature  of  pharmacy,  has  inter- 
fered somewhat  with  the  introduction  of  the  word  in  the  sense  in 
which  it  is  here  defined. 

It  is  fair  to  say,  too,  that  the  distinction  between  hydrates  and 
hydroxides  is  more  or  less  arbitrary,  as  some  hydroxides  lose  water 
more  easily  than  some  hydrates,  and  in  many  cases  the  water  of 
hydration  cannot  be  removed  without  decomposition  of  the  rest 
of  the  salt. 


NATURAL  WATERS  83 

contains  air  in  solution  and  a  little  carbonic  acid  from  the  carbon 
dioxide  of  the  air.  On  falling  upon  the  ground  rain  water  begins 
at  once  to  take  up  various  substances,  partly  in  suspension,  partly 
in  solution.  Calcium  or  magnesium  carbonate  and  calcium 
sulfate  cause  the  water  to  become  hard  (pp.  310,  311)  and  injure 
it  seriously  for  use  in  steam  boilers  or  in  laundries.  If  the  water 
is  mixed  with  sewage,  it  frequently  becomes  contaminated  with 
bacteria  which  produce  disease.  The  diseases  of  typhoid  fever 
and  of  cholera,  especially,  are  frequently  transmitted  in  this  way. 
During  an  epidemic  of  cholera  in  the  cities  of  Hamburg'  and 
Altona,  Germany,  the  people  in  the  houses  on  one  side  of  a  certain 
street  used  a  contaminated  water  which  caused  very  many  cases 
of  the  disease.  Across  the  street,  water  from  the  same  source, 
but  after  passing  a  system  of  public  filters,  was  used  and  there 
were  very  few  cases.  In  Chicago"  before  the  opening  of  the 
drainage  canal  there  were  170  deaths  from  typhoid  fever  per 
year  for  each  100,000  people.  After  the  drainage  canal  carried 
away  the  sewage  which  had  formerly  gone  into  Lake  Michigan 
and  contaminated  the  water  supply  of  the  city,  the  death  rate 
from  typhoid  fell  to  16  per  100,000. 

Purification  of  Water.  Water  may  be  purified  most  com- 
pletely by  distillation,  though  special  precautions  are  required 
to  get  rid  of  ammonia,  carbon  dioxide  and  other  volatile  impuri- 
ties. The  bacteria  in  water  may  be  killed  almost  completely 
by  boiling  the  water  for  a  short  time,  and  this  should  always  be 
done  when  it  is  necessary  to  use  a  suspected  water  for  drinking 
or  the  preparation  of  food.  The  bacteria  may  also  be  almost 
completely  removed  by  filtration,  either  on  a  large  scale  on  beds 
of  sand,  or  through  filters  of  fine-grained  stone  or  unglazed  por- 
celain. Charcoal  filters,  which  were  formerly  used,  are  rarely 
effective.  Waters  may  also  be  sterilized  by  treatment  with 
ozone,  with  ultra-violet  light,  or  with  bleaching  powder. 

Hydrogen  Peroxide.  A  second  compound  of  hydrogen  with 
oxygen,  called  hydrogen  peroxide  and  having  the  formula  H2O2, 
can  be  prepared  by  the  action  of  acids  on  peroxides  of  univalent 
or  bivalent  metals : 


84 


A  TEXTBOOK  OF  CHEMISTRY 


Na2O2         +  2  HC1     =     2  NaCl      +          H2O2 

Sodium  Peroxide  Hydrogen  Peroxide 


Ba02          -f  H2SO4     =     BaSO4 

Barium  Peroxide  Barium  Sulfate 


H2O 


Barium  sulfate  is  almost  insoluble  in  water,  and  if  the  barium 
peroxide  and  sulfuric  acid  are  used  in  equiv- 
alent amounts  a  solution  may  be  obtained 
which  contains  practically  nothing  except 
hydrogen  peroxide  and  water.  Such  a  solu- 
tion may  be  concentrated  in  a  vacuum  des- 
iccator (Fig.  35)  over  sulfuric  acid,  the 
vapor  pressure  of  water  being  much  greater 
than  that  of  hydrogen  peroxide,  or  it  may 
be  distilled  under  the  low  pressure  obtained 
by  means  of  a  good  air  pump  or  filter  pump. 
Water  will  distil  away  first,  and  finally  the 
hydrogen  peroxide  will  distil,  almost  pure. 
Hydrogen  peroxide  is  a  heavy  liquid  which 
decomposes  slowly  at  ordinary  temperatures 


Fig.  35 


into  water  and  oxygen  : 

H2O2  =  H20  +  O 

This  reaction  is  accompanied  by  the  evolution  of  46,200  cal- 
ories of  heat  (by  the  decomposition  of  68  grams  of  hydrogen 
peroxide),  and,  as  is  usual  with  reactions  evolving  heat  and  lib- 
erating a  gas,  may  become  explosive.  The  reaction  is  catalyzed 
by  many  substances,  especially  by  finely  divided  platinum  and 
other  metals,  and  it  is  quite  dangerous  to  bring  concentrated 
solutions  of  hydrogen  peroxide  into  contact  with  organic  sub- 
stances. 

Hydrogen  peroxide  is  a  powerful  oxidizing  agent,  giving  oxygen 
readily  to  many  substances.  This  property,  also,  is  closely  re- 
lated with  the  fact  that  heat  is  evolved  when  it  decomposes. 

If  a  solution  of  hydrogen  peroxide  is  added  to  the  black  pre- 


HYDROGEN  PEROXIDE  85 

cipitate  of  lead  sulfide  suspended  in  water,  the  substance  is  oxi- 
dized to  white  lead  sulfate  : 

PbS  +  4  H2O2  =  PbSO4  +  4  H2O 

Hydrogen  peroxide  also  acts  on  silver  oxide  and  on  a  good  many 
other  substances  as  a  reducing  agent,  but  in  all  such  cases  oxy- 
gen is  liberated.  We  may  explain  this  by  considering  that 
oxygen  gas  is  a  compound  of  oxygen  with  itself l  and  that  the  oxy- 
gen of  the  hydrogen  peroxide  oxidizes  the  oxygen  of  the  silver 
oxide  to  free  oxygen.  Or  it  may  be  that  the  oxygen  of  the 
silver  oxide  oxidizes  the  hydrogen  of  the  hydrogen  peroxide  to 
water,  leaving  the  two  oxygen  atoms  of  the  peroxide  combined 
together  as  free  oxygen,  O22 : 

H202  +  Ag20  =  H20  +  02  +  2Ag 

The  reduction  of  potassium  permanganate,  KMn04,  in 
acid  solution  is  frequently  used  for  the  quantitative  deter- 
mination of  hydrogen  peroxide.  The  equation  representing 
this  reaction  is  most  easily  written  in  two  parts,  as  indicated 
below.  The  first  part  of  the  reaction  is  based  on  the  facts  that 
potassium  is  univalent  and  manganese  bivalent  when  combined 
with  acid  radicals  and  that,  for  this  reason,  when  two  molecules 
of  potassium  permanganate  are  acted  on  by  sulfuric  acid  in  the 
presence  of  some  substance  which  can  take  up  oxygen,  five  atoms 
of  oxygen  become  available  for  the  oxidation  of  that  other  sub- 
stance. These  five  atoms  of  oxygen  are  inclosed  in  brackets 
to  indicate  that  the  decomposition  does  not  take  place  except  in 
the  presence  of  something  with  which  this  oxygen  can  combine : 

2  KMn04  +  3  H2SO4  =  K2SO4  +  2  MnSO4  +  3  H2O  +  (5  O) 

(5  O)  +  5  H2O2  =  5  H2O  +  5  O2 
2  KMnO4  +  3  H2SO4  +  5  H2O2 

=  K2SO4  +  2  MnSO4  +  8  H2O  +  5  O2 

1  Other  reasons  for  such  a  view  will  be  given  later  (p.  93). 

2  The  fact  that  the  reaction 

H2O2  =  H2O  +  O 
is  monomolecular  points  strongly  to  the  latter  explanation. 


86  A  TEXTBOOK  OF  CHEMISTRY 

The  last  equation  is  obtained  by  combining  the  other  two 
algebraically,  eliminating  the  five  atoms  of  oxygen  which  appear 
on  opposite  sides. 

*  Properties   and  Uses   of  Hydrogen  Peroxide.  ~  Hydrogen 
peroxide  has  a  specific  gravity  of  1.4584  at  0°.     It  boils  at  69.2° 
under  a  pressure  of  26  mm.  or  at  84°-85°  under  a  pressure  of 
68  mm.     A  dilute  solution  is  fairly  stable  when  pure,  the  stabil- 
ity being  increased  by  the  presence  of  a  small  amount  of  a  mineral 
acid.     The  stability  is  decreased  by  alkalies  and  by  many  other 
substances,  and  the  decomposition  is  also  hastened  by  exposure 
to  light. 

Hydrogen  peroxide  is  used  in  medicine  as  a  bactericide  and  for 
the  diagnosis  of  pus,  which  causes  its  rapid  decomposition  with 
evolution  of  oxygen.  The  medicinal  solution  commonly  used 
is  known  as  a  10-volume  solution,  meaning  that  it  evolves  ten 
times  its  volume  of  oxygen  when  it  decomposes,  or  20  times  its 
volume  when  treated  with  an  oxidizing  agent,  one  half  of  the 
oxygen  coming  from  the  latter,  as  explained  above. 

Hydrogen  peroxide  is  also  used  to  bleach  hair,  silk  and  wool, 
being  much  more  suitable  than  chlorine  or  hypochlorites  (p.  127) 
for  this  purpose.  A  solution  suitable  for  this  purpose  may  be 
obtained  by  dissolving  sodium  peroxide  in  cold  water  and  adding 
dilute  sulfuric  acid. 

*  Tests  for  Hydrogen  Peroxide.     With  potassium  dichromate 
in  an  acid  solution  hydrogen  peroxide  gives  a  beautiful  blue 
compound  which  is  soluble  in  ether.     The  composition  of  the 
compound  is  not   positively  known.      Its   formation   may  be 
used  either  as  a  test  for  hydrogen  peroxide  or  as  a  test  for  chro- 
mium.    Another  valuable  test  is  the  yellow  color  given  with 
solutions  containing  titanium. 

*  Structure    of   Hydrogen   Peroxide.      Two   formulas   have 
been  proposed  to  represent  the  structure  of  hydrogen  peroxide : 


H— O— O— H  and      "€> = O 


LAW  OF  MULTIPLE  PROPORTION  87 

The  first  represents  both  oxygen  atoms  as  bivalent,  the  second 
represents  one  of  the  oxygen  atoms  as  quadrivalent.  The  evi- 
dence in  favor  of  the  first  formula  is  : 

1.  There  are  a  number  of  reactions  in  which  hydrogen  peroxide 
seems  to  be  formed  by  the  reduction  of  oxygen,  which  has  the 
formula  O2  (p.  93)  : 

O=O  +  H-H  =  H—  O—  O—  H 

2.  Hydrogen  peroxide  does  not  seem  to  be  formed  by  the  oxi- 
dation of  water,  as  it  should  be  if  the  second  formula  were  true. 

3.  When  the  two  hydrogen  atoms  of  hydrogen  peroxide  are 
replaced  by  the  ethyl  group,  C2H5,  diethyl  peroxide,  (€2115)202, 
is  formed  and  this  gives  ethyl  alcohol,  C2H5  —  O  —  H,  by  re- 
duction.    This  is  easily  explained  by  the  first  formula  : 

OjjHs  —  o  —  o  —  G2ri5  -p  ±12  ==  2  C^HS  —  o  —  M 

If  diethyl  peroxide  had  a  formula  corresponding  to  the 
second  one  for  hydrogen  peroxide,  it  should  give  ethyl  ether, 
C2H5  —  O  —  C2H5  by  reduction  (Baeyer  and  Villiger,  Ber.  33, 
3387  (1900))  : 


C2H5 

Law  of  Multiple  Proportion.  In  water  and  hydrogen  peroxide 
the  same  elements,  hydrogen  and  oxygen,  combine  to  form  two 
different  compounds.  In  water  one  part  of  hydrogen  combines 
with  eight  parts  of  oxygen  (or  2  with  16),  while  in  hydrogen  per- 
oxide one  part  of  hydrogen  combines  with  sixteen  parts  of  oxy- 
gen (or  2  with  32).  Nitrogen  and  oxygen  form  a  whole  series 
of  compounds  represented  by  the  formulas  and  composition  : 

N:O 

Nitrous  oxide,  N2O  28  :  16 

Nitric  oxide,  NO  14  :  16 

Nitrogen  trioxide,  N2O3  28  :  48 
Nitrogen  tetroxide,  N2O4  28  :  64 
Nitrogen  pentoxide,  N2O5  28  :  80 


88  A  TEXTBOOK  OF  CHEMISTRY 

In  every  such  case,  if  we  consider  some  fixed  amount  of  one  oj 
the  elements  (it  makes  no  difference  which  one]  the  amounts  of  the 
other  combining  with  this  fixed  amount  will  bear  a  simple  ratio  to 
each  other.  This  is  known  as  the  Law  of  multiple  proportions. 
The  discovery  of  this  law  led  Dalton  to  propose  the  atomic 
theory  of  the  constitution  of  matter.  A  little  consideration  of 
the  law  shows  that  it  follows,  necessarily,  from  the  law  of  com- 
bining weights  (p.  13),  and  that  the  law  of  combining  weights  is 
more  comprehensive  and  important. 

Two  compounds  of  iron  contain : 

Ferrous  oxide,    77.73  per  cent  of  iron 

22.27  per  cent  of  oxygen 
100. 

Ferric  oxide,      69.94  per  cent  of  iron 

30.06  per  cent  of  oxygen 

To67~ 

Show  that  these  proportions  are  in  accordance  with  the  law 
of  multiple  proportion.  This  illustration  helps  us  to  understand 
why  the  law  was  not  discovered  before  the  nineteenth  century. 


CHAPTER  VI 


AVOGADRO'S  LAW. 


SELECTION  OF  ATOMIC  WEIGHTS. 
OZONE 


Gay  Lussac's  Law  of  Combining  Volumes.  It  has  been  shown 
that  hydrogen  and  oxygen  unite  very  nearly  in  the  proportion 
of  two  to  one  by  volume  and  that  the  volume  of  steam  formed  is 
very  nearly  the  same  as  the  volume  of  hydrogen  which  it  contains. 
A  study  of  many  other  gaseous  elements  and  their  compounds 
has  shown  that  in  every  case  there  is  a  simple  ratio  between  the 
volumes  of  gases  which  combine  with  each  other  and  also  between 
those  volumes  and  the  volume  of  the  product,  if  that  is  a  gas.  This 
is  known  as  Gay  Lussac's  law  of  combining  volumes,  and  it  is 
true  of  all  elements  or  compounds  which  can  be  converted  into 
gases  without  decomposition,  as  well  as  of  substances  which  are 
gases  at  ordinary  temperatures.  It  has  been  shown  to  be  true 
for  thousands  of  compounds.  The  law  may  be  illustrated  by  the 
following  diagrams : 


36.5  grains 

36.5  grams 

Hydrocloric 

Hydrocloric 

acid 

acid 

90 


A  TEXTBOOK  OF  CHEMISTRY 


2  grams 
Hydrogen 


2  grams 
Hydrogen 


2  grams 
Hydrogen 


17  grams 
Ammonia 

17  grams 
Ammonia 

2  grams 

2  grams 

Hydrogen 

Hydrogen 

2  grams 

2  grams 

Hydrogen 

Hydrogen 

2  grams 

2  grams 

Hydrogen 

Hydrogen 

34  grams 

34  grams 

Phosphine 

Phosphine 

34  grams 

34  grams 

Phosphine 

Phosphine 

The  volumes  have  been  chosen  for  these  illustrations  in  such  a 
way  that  the  unit  volume  always  contains  one  gram,  or  a  whole 
number  of  grams,  of  hydrogen.  In  every  case  the  weight  of  this 
unit  volume  bears  a  simple  relation  to  the  atomic  weights  of  the  ele- 
ments which  it  contains.  This  volume  of  any  gaseous  element  or 
compound  always  contains  one  gram-atom  or  a  whole  number  of 
gram-atoms  of  each  element  contained  in  the  gas.  This  is  not 
accidental,  but  follows  of  necessity  from  the  two  laws :  (V\  that 


AVOGADRO'S  LAW  91 

the  composition  of  every  compound  can  be  expressed  by  multi- 
ples of  the  atomic  weights  of  the  elements  which  compose  it ; 
and  (2)  that  there  is  always  a  simple  ratio  between  the  volumes 
of  gases  which  combine  and  also  between  those  volumes  and  the 
volume  of  the  product,  if  that  is  a  gas. 

A  further  examination  of  the  illustrations  shows  that  the  unit 
volume  of  each  compound  contains  one  gram  molecule  of  the 
compound.  This  result  does  not  follow  of  necessity  from  the 
two  laws  just  given.  It  depends  on  the  values  which  we  give  to 
the  atomic  weights.  Thus  if  we  were  to  call  the  atomic  weight 
of  oxygen  8  and  the  formula  of  water  HO,  as  was  done  by  Dalton, 
the  gram  molecular  weight  of  water  would  be  9  and  the  unit 
volume  which  has  been  chosen  would  contain  two  gram  mole- 
cules of  water.  As  has  been  stated,  we  have  chosen  as  the  unit 
volume  for  our  illustration  that  volume  which,  in  several  com- 
pounds, contains  one  gram  of  hydrogen,  of  course  because  hy- 
drogen is  our  (approximate)  unit  for  atomic  weights.  No  com- 
pound of  oxygen  is  known  which  contains  less  than  16  grams  in 
this  unit  volume. 

Avogadro's  Law.  The  selection  of  16  instead  of  8  as  the  atomic 
weight  of  oxygen  is  based  on  a  hypothesis  proposed  by  Avogadro, 
an  Italian  chemist,  in  1811.  This  hypothesis  is  that  all  gases, 
under  the  same  conditions  of  temperature  and  pressure,  contain 
equal  numbers  of  molecules  in  equal  volumes.  Some  of  the  many 
facts  which  support  this  hypothesis  so  fully  that  it  may  now  be 
considered  as  an  established  law  are  : 

1.  While  solids  and  liquids  vary  greatly  in  their  rate  of  expan- 
sion or  contraction  for  changes  of  pressure  or  of  temperature,  all 
gases  expand  and  contract  alike.     This  points  very  strongly  to 
a  similarity  in  their  structure. 

2.  When  the  law  of  combining  weights  is  combined  with  the 
law  of  combining  volumes  it  follows,  of  necessity,  that  there  must 
be  a  simple  ratio  between  the  numbers  of  molecules  in  equal 
volumes  of  different  gases.      But  if  the  ratio  is  one  of  simple 
whole  numbers,  it  seems  highly  improbable  that  it  is  not  one  of 
equality.     Thus  it  would  seem  very  improbable  that  the  number 


92  A  TEXTBOOK  OF  CHEMISTRY 

of  molecules  in  a  given  volume  of  steam  is  exactly  twice  the 
number  in  the  same  volume  of  hydrochloric  acid,  as  we  should 
have  to  suppose  if  the  molecular  weight  of  water  is  9  while  that 
of  hydrochloric  acid  is  36.5. 

3.  The  kinetic  theory  of  gases,  which  accounts  so  well  for  their 
properties,  leads  directly  to  Avogadro's  law  on  the  basis  of  the 
laws  of  the  collision  of  elastic  bodies  (p.  58). 

4.  The  atomic  weights  selected  on  the  basis  of  Avogadro's 
law  have  made  possible  the  classification  of   the  elements  in 
groups  according  to  their  atomic  weights.     This  classification 
is  known  as  the  Periodic  System  (p.  132)  and  furnishes  very 
strong,  independent  evidence  that  the  atomic  weights  selected 
are  in  reality  the  true  relative  weights  of  the  atoms. 

Selection  of  an  Atomic  Weight.  According  to  Avogadro's  law, 
equal  volumes  of  different  gases,  under  the  same  conditions  of 
temperature  and  pressure,  must  have  weights  proportional  to 
the  weights  of  the  molecules  of  the  gases.  If  the  molecule  of  one 
gas  is  twice  as  heavy  as  that  of  another,  one  liter  of  the  first  gas 
must  weigh  twice  as  much  as  a  liter  of  the  second.  Accordingly 
if  we  can  take  as  our  unit  volume  the  volume  filled  by  a  gram 
molecule  of  some  compound  which  contains  only  one  atom  of 
hydrogen  in  its  molecule,  this  unit  volume  will  contain  one  gram 
molecule  of  every  other  gaseous  element  or  compound.  We  can- 
not, of  course,  look  at  the  molecules  of  different  compounds  to 
discover  which  one  contains  only  a  single  atom  of  hydrogen ;  but 
if  we  weigh  the  same  volume  of  hydrochloric  acid,  steam  and 
ammonia,  we  find  that  in  a  given  volume  of  steam  there  is  twice 
as  much  hydrogen  as  there  is  in  the  same  volume  of  hydrochloric 
acid,  and  that  in  the  same  volume  of  ammonia  there  is  three  times 
as  much  hydrogen.  As  no  compound  has  ever  been  found  which 
contains  less  hydrogen  than  hydrochloric  acid  does  in  the  unit 
volume,  it  seems  pretty  safe  to  conclude  that  there  is  only  one 
atom  of  hydrogen  in  a  molecule  of  this  compound.  As  one  gram 
of  hydrogen  combines  with  35.5  grams  of  chlorine,1  the  volume 

1  Approximate  values  are  used  for  convenience,  as  always.  The 
true  values  are  1.0078  grams  of  hydrogen  for  35.46  grams  of  chlorine. 


SELECTION  OF  ATOMIC  WEIG-HTS 


93 


filled  by  36.5  grams  of  hydrochloric  acid  will  contain  one  gram 
molecule,  and  it  must  contain  one  gram  molecule  of  every  other 
gas. 

In  the  discussion  above  it  is  pointed  out  that  those  compounds 
of  hydrogen  which  contain  the  smallest  amount  of  hydrogen  in 
the  unit  volume  probably  contain  only  one  atom  of  hydrogen 
in  their  molecules.  Reasoning  in  the  same  way,  we  may  find 
for  each  element  those  compounds  which  contain  the  smallest 
amount  of  the  element  in  the  unit  volume,  and  it  is  probable  that 
these  compounds  contain  only  one  atom  of  the  element  in  the 
molecule.  Thus  steam  contains  16  grams  of  oxygen  in  the 
unit  volume ;  and  as  no  compound  of  oxygen  containing  a  smaller 
amount  is  known,  we  conclude  that  there  is  only  one  atom  of 
oxygen  in  a  molecule  of  steam  and  that  the  atomic  weight  of 
oxygen  is  16.  The  atomic  weights  of  the  other  elements  which 
form  gaseous  compounds,  or  compounds  which  can  be  converted 
into  gases  without  decomposition,  have  been  selected  in  the  same 
way. 

Molecules  of  the  Elements.  It  was  natural  for  Dalton  when 
he  proposed  the  atomic  theory  to  think  of  the  atom  as  the  small- 
est particle  of  an  element  in  the  free  state,  and  it  did  not  occur 
to  him  that  atoms  of  the  same  kind  could  combine.  A  refer- 
ence to  the  diagrams,  however,  shows  that  there  is  twice  as  much 
oxygen  in  the  unit  volume  of  oxygen  gas  as  in  the  unit  volume 
of  steam.  According  to  Avogadro's  law  it  follows  that  oxygen 
gas  contains  two  atoms  of  oxygen  in  each  molecule.  We  may 
reach  the  same  conclusion  by  another  process.  If  we  let  each 
square  below  represent  1000  molecules,  it  is  clear  that  2000 
molecules  of  steam  are  formed  from  1000  molecules  of  oxygen, 
and,  as  each  molecule  of  steam  must  contain  at  least  one  atom 


94  A  TEXTBOOK  OF  CHEMISTRY 

of  oxygen,  the  1000  molecules  of  oxygen  must  consist  of  2000 
atoms,  or  each  molecule  contains  two  atoms. 

It  will  be  seen  from  the  diagrams  that  four  atoms  combine 
to  form  a  molecule  of  phosphorus  vapor,  while  a  molecule  of 
mercury  vapor  contains  only  a  single  atom.  In  the  latter  case 
atom  and  molecule  are  identical.  In  general  the  atoms  of  non- 
metallic  elements  combine  to  form  molecules  of  the  element  in 
the  free  state,  but  the  atoms  of  the  metals  show  little  tendency 
to  combine  in  this  way.  For  a  probable  explanation  of  this  re- 
markable difference  see  J.  J.  Thompson,  The  Corpuscular  Theory 
of  Matter,  p.  120  ;  and  H.  N.  McCoy,  J.  Am.  Chem.  Soc.  33,  273. 

Gram  Molecular  Volume.  The  unit  volume,  which  will  con- 
tain one  gram  molecule  of  any  gas,  is  best  calculated  from  the 
weight  of  a  liter  of  oxygen,  since  oxygen  is  the  basis  for  atomic 
weights  (p.  68).  One  gram  molecule  of  oxygen,  C>2,  contains 
32  grams,  and,  as  a  liter  of  oxygen  weighs  1.429  grams,  the  gram 

32    __ 
molecular  volume  for  oxygen  will  be  =  22.4  liters  at  0° 


and  760  mm.  At  the  same  temperature  and  pressure  one 
gram  molecule  of  any  other  gas  will  fill  the  same  volume,  22.4 
liters. 

This  statement  and  the  law  of  Avogadro  are  subject  to  limita- 
tions similar  to  those  which  apply  to  the  laws  of  Boyle  and 
Charles.  Just  as  most  gases,  and  especially  those  which  are 
easily  liquefied,  contract  too  much  when  compressed  from  one 
to  two  atmospheres  pressure  and  also  contract  too  much  when 
cooled  from  100°  to  0°,  so  almost  all  gases  are  heavier  than  they 
should  be  in  accordance  with  the  law  of  Avogadro.  As  the 
volume  increases  under  diminished  pressure,  however,  gases  ap- 
proach the  condition  of  an  "ideal  "  gas,  and  at  low  pressures  the 
densities  of  gases  agree  very  closely  indeed  with  the  law.  The 
amount  of  the  deviation  from  the  law  at  atmospheric  pressure 
and  the  agreement  under  low  pressures  will  be  obvious  from  the 
following  table  : 


DENSITY  OF  GASES 


95 


DENSITY  OF  GASES  l 


NAME 

FOBMULA 

WEIGHT  OF 
ONE  LITER 

WEIGHT  OF 
22.4  LITERS 

AT   0°  AND 
760  MM. 

DENSITY  AT 
Low 
PRESSURE 
0  =  32 

MOLECU- 
LAR 
WEIGHT 

Oxygen      .     .     . 

02 

1.429 

32.00 

32.00 

32. 

Hydrogen       .     . 

H2 

0.08987 

1.997 

2.01 

2.016 

Nitrogen    .     .     . 

N2 

1.2507 

28.02 

28.01 

28.02 

Carbon  monoxide 

CO 

1.2504 

28.01 

28.00 

28.00 

Nitric  oxide   .     . 

NO 

1.3402 

30.02 

30.01 

30.01 

Argon   .... 

Ar 

1.7808 

39.89 

39.88 

39.88 

Carbon  dioxide  . 

C02 

1.9768 

44.28 

44.01 

44.00 

Nitrous  oxide 

N20 

1.9777 

44.30 

44.03 

44.02 

Hydrochloric  acid 

HC1 

1.6398 

36.73 

36.47 

36.47 

Ammonia       .     . 

NH3 

0.7708 

17.27 

17.01 

17.03 

Sulfur  dioxide     . 

SO2 

2.9266 

65.56 

64.07 

64.07 

Air   

1.2928 

28.96 



The  table  shows  that  at  low  pressures  the  deviations  from  Avo- 
gadro's  law  scarcely  exceed  the  experimental  errors  of  the  deter- 
minations. If  the  pressure  and  volume  of  a  gas  are  corrected  for 
the  volumes  occupied  by  the  molecules  and  for  their  attractions  for 
each  other,  corrections  which  can  be  determined  experimentally 
(van  der  Waals) ,  the  law  of  Avogadro  also  becomes  almost  rigor- 
ously exact. 

The  weight  of  22.4  liters  of  air  furnishes  a  very  simple  method 
of  calculating  the  density  of  any  gas  as  compared  with  air. 
For  approximate  estimates  the  weight  may  be  taken  as  29  grams, 
and  the  molecular  weight  of  any  compound  divided  by  29  will 
be  approximately  its  density  as  compared '  with  that  of  air. 
Thus  it  will  be  seen  that  hydrogen  is  14 J  times  lighter  than  air, 
while  oxygen  is  1.1  times  as  heavy  and  carbon  dioxide  one  and 
one  half  times  as  heavy. 

*  Number  of  Molecules  in  one  Cubic  Centimeter  of  a  Gas. 
Avogadro's  law  was  established  with  a  high  degree  of  probability 
before  any  means  of  estimating  the  number  of  molecules  in  a 


1  Guye,  J.  Am.  Chem.  Soc.  30,  155  (1908). 


96  A  TEXTBOOK  OF  CHEMISTRY 

given  volume  of  a  gas  was  discovered.  It  is  of  some  interest 
to  know,  however,  that  several  different  methods  of  estimating 
this  number  are  now  known  and  that  the  results  obtained  by 
different  methods  are  in  fair  agreement.  Two  of  these  methods 
may  be  given  here,  in  outline. 

It  has  been  shown  that  the  element  radium  slowly  decomposes 
and  that  as  it  does  so  it  shoots  out  atoms  of  helium  with  a  tremen- 
dous velocity.  The  volume  of  helium  given  off  by  a  gram  of 
radium  has  been  measured  and  is  0.46  cubic  millimeter  in  a  day. 
When  an  atom  of  helium  shot  out  by  radium  hits  a  screen  of  zinc 
sulfide,  it  produces  a  flash  of  light ;  and  as  the  helium  atoms  are 
sent  out  in  all  directions  equally,  by  placing  a  screen  of  zinc 
sulfide  back  of  a  small  opening  which  is  at  a  known  distance 
from  a  weighed  amount  of  radium  and  counting  the  flashes,  it 
is  possible  to  estimate  the  number  of  atoms  of  helium  which  pass 
the  opening  in  a  given  time  and  so  the  number  of  atoms  in  a 
cubic  centimeter  of  the  gas.  Rutherford  has  estimated  the 
number  of  molecules  in  one  cubic  centimeter  of  helium,  in  this 
way,  as  2.56  X  1019  (or  25,600,000,000,000,000,000).  (Report  of 
the  Winnipeg  Meeting  of  the  British  Association,  1909, 
p.  377.) 

When  very  minute  particles  suspended  in  water  are  observed 
with  a  microscope,  it  is  seen  that  they  are  never  at  rest,  but  move 
about  constantly  in  a  wholly  irregular  manner.  This  was  first 
noticed  by  the  English  botanist  Brown  in  1827  and  is  called  the 
Brownian  movement.  It  has  long  been  considered  as  an  evi- 
dence that  water  and  other  liquids  are  composed  of  molecules 
which  are  in  rapid  motion.  Perrin  has  succeeded  in  showing 
(1909)  that  in  an  emulsion  of  gamboge  the  minute  particles  are 
distributed  as  they  should  be  if  they  are  considered  as  very  large 
molecules,  and  he  has  connected  this  distribution  with  the  kinetic 
theory  in  such  a  manner  as  to  furnish  an  estimate  of  the  number 
of  molecules  in  a  cubic  centimeter  of  a  gas,  which  he  gives  as 
3.15  X  1019.  A  number  of  other  methods  of  estimating  the 
same  quantity  give  results  of  the  same  general  value.  As  the 
different  methods  are  quite  independent  of  each  other,  we  can 


OZONE 


97 


have  considerable  confidence  that  the  values  are  approximately 
correct.1 

Allotropic  Forms.  Ozone.  When  oxygen  is  subjected  to  the 
action  of  electrical  waves  obtained  by  connecting  the  tinfoil  at 
A  and  D  (Fig.  36),  with  the  poles  of  an  induction  coil,  from  5 
to  8  per  cent  of  the  gas  is  converted  into  ozone.  If  the  mixture 


•*• 

Q 
1 

r^A-                                             (( 

w\ 

^ 

1 

1 

B 

% 

=:i                                       D  : 

I 

^•J. 

1 

===^ 

fa 

1 

112 

Fig.  36 

of  oxygen  and  ozone  is  partly  liquefied  by  passing  it  into  a  flask 
immersed  in  liquid  air,  a  dark  blue  liquid  is  obtained.  Oxygen 
boils  at  —  182.5°,  while  ozone  boils  at  —  119°,  and  a  mixture  of 
oxygen  and  ozone  containing  84  per  cent  of  ozone  has  been  ob- 
tained by  allowing  the  oxygen  to  boil  away  from  such  a  mixture. 
(Ladenburg,  Ber.  31,  2508,  2830  (1898).) 

Ozone  has  a  strong  odor,  which  is  noticed  in  the  neighborhood 
of  electrical  machines.  The  weight  of  the  gram  molecular 
volume  (22.4  liters)  of  the  gas  is  48  grams.  Hence  the  formula 
is  O3.  Ozone  is  evidently  formed  in  accordance  with  the  equa- 
tion : 


The  reaction  is  reversible  and  the  ozone  formed  is  unstable, 
decomposing  slowly  at  ordinary  temperatures,  rapidly  and 
completely  at  250°-300°.  The  decomposition  is  accompanied  by 
considerable  evolution  of  heat,  and  so  liquid  ozone  may  easily 
give  violent  explosions.  One  gram  molecule,  48  grams,  gives 

1  Professor  R.  A.  Millikan  of  the  University  of  Chicago  has 
recently  determined  the  number  by  a  new  method  which  is,  appar- 
ently, much  more  accurate  than  any  previously  used.  He  gives 
the  value  2.71  x  1019. 


98  A  TEXTBOOK  OF  CHEMISTRY 

out  29,400  calories  in  decomposing,  and  the  properties  of  ozone 
are  intimately  connected  with  the  fact  that  it  contains  so  much 
more  chemical  energy  than  ordinary  oxygen.  It  is  a  vigorous 
oxidizing  agent  and  attacks  metallic  silver  and  many  other  sub- 
stances which  are  not  affected  by  ordinary  oxygen. 

Ozone  may  be  detected  by  its  action  on  moist  potassium 
iodide  starch  paper : 

O3  -f  2  KI  +  H2O  =  2  KOH  +  I2  -f  O2 

The  liberated  iodine  gives  a  deep  blue  color  with  the  starch. 
A  small  amount  of  ozone  is  formed  when  a  piece  of  clean  phos- 
phorus, half  covered  with  water,  is  allowed  to  stand  for  a  short 
time  in  a  bottle  rilled  with  air.  The  ozone  may  be  detected  by 
the  odor  and  by  the  potassium  iodide  starch  paper. 

Ozone  has  been  used  to  a  limited  extent  for  sterilizing  drinking 
water,  but  this  use  has  not  proved  very  successful.  Bleaching 
powder  is  cheaper  and  much  more  efficient.  Ozone  is,  however, 
a  powerful  germicide,  and  it  may  be  that  the  ozone  generated  in 
thunderstorms  plays  a  beneficent  part  in  nature. 

In  ordinary  oxygen  and  ozone  we  have  two  different  forms  of 
the  same  element,  and  it  is  evident  that  in  this  case  we  may  ob- 
tain another  substance  from  an  element  without  adding  any- 
thing to  it.  The  formulas  of  the  two  forms  of  oxygen  give  us  a 
partial  explanation  of  the  difference  between  the  two  substances. 
Both  forms  are  really  compounds  in  which  oxygen  is  combined 
with  itself.  We  can  represent  this  graphically  by  the  formulas  : 

O 
O=O  and  /\  or  O  =  O=O 

o— o 

The  last  formula  represents  one  of  the  oxygen  atoms  as 
quadrivalent. 

Several  elements  beside  oxygen  exist  in  two  or  more  forms. 
Such  forms  are  called  allotropic. 


OZONE  99 

EXERCISES 

1.  What  will  be  the  volume  of  a  gram  molecule  of  oxygen  at  20° 
and  750  mm.  ? 

2.  How  many  grams  of  sodium  peroxide  will  be  required  to  give 
22.4  liters  of  oxygen  under  standard  conditions  ? 

3.  How  many  grams  of  potassium  chlorate  will  be  required  to  give 
22.4  liters  of  oxygen  ? 

4.  How  many  liters  of  carbon  dioxide  will  be  formed  by  burning  6 
grams  of  carbon  ? 

5.  How  many  liters  of  carbon  monoxide  ?     How  many  liters  of 
oxygen  will  be  required  for  the  reaction  in  each  case  ? 

6.  How  many  grams  of  iron  will  be  required  to  give  22.4  liters  of 
hydrogen,  if  the  iron  is  dissolved  in  hydrochloric  acid  ?     How  many 
grams,  if  dissolved  in  sulfuric  acid-?     How  many  grams,  if  used  to 
decompose  steam  ? 

7.  How  much  heat  could  be  obtained,  if  two  grams  hydrogen  could 
be  burned  in  ozone  ? 

8.  How  many  grams  of  sodium  are  required  to  give  22.4  liters  of 
hydrogen  ? 

9.  How  many  milligrams  of  aluminium  are  required  to  give  22.4  cc. 
of  hydrogen  ? 


CHAPTER  VII 

CHLORINE 

SYMBOL,  Cl.    ATOMIC  WEIGHT,  35.46.    FORMULA,  C12. 

Occurrence  of  Chlorine.  Chlorine  is  not  found  free  in  nature, 
chiefly  because  of  its  strong  affinity  for  almost  every  other  ele- 
ment and  especially  for  the  metals.  Its  most  important  com- 
pound is  common  salt,  sodium  chloride,  NaCl.  This  is  found  in 
large  amounts  in  sea  water,  in  strong  brines  from  artesian  wells 
in  very  many  places  and  in  enormous  beds  of  rock  salt,  some- 
times a  hundred  feet  in  thickness.  Calcium  chloride  and  mag- 
nesium chloride  are  also  found  in  sea  water  and  in  many  of  the 
brines.  Silver  chloride  is  sometimes  an  important  ore  of  silver. 

Preparation  of  Chlorine.  1.  By  Electrolysis  of  Sodium  Chlo- 
ride. If  an  electric  current  is  passed  through  a  solution  of 
common  salt,  NaCl,  using  a  carbon  anode  and  mercury  as  a 
cathode,  the  negative  chloride  ions,  Cl~,  will  be  discharged  at  the 
anode  and  the  chlorine  will  be  evolved  as  a  gas.  The  sodium 
will  dissolve  in  the  mercury  cathode,  and  by  appropriate  means  it 
may  be  caused  to  react  with  water,  giving  sodium  hydroxide  and 
hydrogen.  This  process  is  now  used  extensively  in  the  manu- 
facture of  chlorine  and  caustic  soda  or  sodium  hydroxide.  It 
is  known  as  the  Castner-Kellner  process  and  will  be  considered 
further  in  connection  with  sodium  hydroxide  (p.  402). 

2.  Preparation  by  Oxidation  of  Hydrochloric  Acid.  In  the 
preparation  of  oxygen,  compounds  (mercuric  oxide  or  potassium 
chlorate)  are  selected  which  decompose  with  liberation  of  the 
gas  when  heated.  In  the  preparation  of  hydrogen,  compounds 
(water  or  hydrochloric  or  sulfuric  acid)  are  selected  from  which 
the  hydrogen  can  be  displaced  by  another  element.  Both 
methods  may  be  used  for  the  preparation  of  chlorine. 

100 


CHLORINE 


101 


In  all  of  the  methods  practically  used  except  electrolysis,  the 
chlorine  of  hydrochloric  acid  is  displaced  by  oxygen,  one  atom 
of  the  bivalent  oxygen  displacing  two  atoms  of  chlorine : 

-H-     ^i    \    /-\  -H-  \/^    I    /"<i      /~<] 

TJ         f^l       I"    ^    —     TT  S^*   T    ^"^         ^l 
±1 d  ll/ 

In  several  methods  compounds  of  manganese  are  used,  and 
these  methods  depend  upon  the  fact  that  while  manganese  forms 
compounds  with  oxygen  in  which  it  is  quadrivalent,  or  in  which 
it  even  shows  a  higher  valence,  the  compounds  of  the  element 
with  chlorine  which  contain  more  than  two  atoms  of  chlorine 
for  one  of  manganese  are  very  unstable.  The  compounds  with 

.0 
oxygen  are  manganese  dioxide,  Mn^  ,  and  potassium  perman- 


ganate,  K — O — Mn^O.     The  stable  compound  with  chlorine 
%0 

is  Mn<g- 

If  a  concentrated  solution  of  hydrochloric  acid  is  poured  on 
manganese  dioxide,  there  is  formed,  at  first,  a  dark  brown  solu- 
tion which  probably  contains  some  manganese  tetrachloride : 

Mn02       +        4HC1       =     MnCl4     +       2  H2O 

Manganese  Hydrochloric          Manganese 

Dioxide  Acid  Tetrachloride 

This  equation  represents  the  reaction  as  a  metathesis  in  which 
two  compounds  each  separate  into  two  parts,  and  then  one  part 
of  each  combines  with  one  part  of  the  other  compound.  A  very 
large  majority  of  chemical  reactions  belong  to  this  class,  and 
the  principle  is  very  useful  in  writing  and  understanding  chemical 
equations.  In  the  present  case  the  manganese  tetrachloride 
is  so  unstable  that  only  indirect  evidence  of  its  existence  has 
been  obtained.  When  the  solution  is  warmed  gently,  it  is  de- 
composed completely  into  manganous  chloride  and  chlorine : 

MnCl4  =  MnCl2  +  C18 


102 


A  TEXTBOOK  '  OF  CHEMISTRY 


The  two  reactions,  which  go  on  simultaneously,  may  be  ex- 
pressed in  one  equation,  thus  : 

Mn02  +  4  HC1  =  MnCl2  +  C12  -f  2  H2O 

If  concentrated  hydrochloric  acid  is  al- 
lowed to  drop  on  potassium  permanganate, 
KMnC>4,  which  is  a  powerful  oxidizing 
agent,  the  hydrochloric  acid  is  oxidized 
and  chlorine  is  liberated.  A  part  of  the 
chlorine,  of  course,  remains  combined 
with  the  potassium  and  manganese.  In 
writing  the  equation  we  notice  that  we 
must  have  8  molecules  of  hydrochloric 
acid  to  furnish  the  hydrogen  to  combine 
with  the  4  atoms  of  oxygen  in  one  mole- 
cule of  the  potassium  permanganate,  and 
we  write,  at  first : 

KMnO4  +  8  HC1  = 

KC1  +  MnCl2  +  4  H2O  +  5  Cl 

But  this  form  of  the  equation  gives  an 
odd  number  of  atoms  of  chlorine,  and  as 
the  chlorine  is  liberated  in  the  molecular 

form  it  is  necessary  to  double  the  equation  to  represent  the 

substances  actually  formed : 

2  KMnO4  +  16  HC1  =  2  KC1  +  2  MnCl2  +  8  H2O  +  5  C12 

3.  Preparation  of  Chlorine  by  the  Deacon  Process.  In  the 
methods  described  in  the  last  paragraph  for  the  preparation  of 
chlorine  the  element  is  liberated  from  hydrochloric  acid  by 
means  of  an  expensive  oxidizing  agent.  Henry  Deacon  of  Eng- 
land devised  a  process  a  good  many  years  ago  by  which  the  ex- 
pensive oxidizing  agent  is  replaced  by  the  oxygen  of  the  air. 
If  a  mixture  of  hydrochloric  acid  and  oxygen,  or  air,  is  heated, 
the  reversible  reaction  represented  by  the  equation  : 

4  HC1  +  O2  ^t  2  H2O  +  2  C12 


Fig.  37 


CHLORINE  103 

takes  place,  but  at  a  temperature  of  350°  to  400°  the  reaction  is 
too  slow  to  be  commercially  possible,  as  time  is  a  very  important 
element  in  all  technical  processes.  The  reaction  might  be  has- 
tened by  using  a  higher  temperature,  but  at  a  high  temperature 
the  equilibrium  between  the  four  substances  is  displaced  in  such 
a  way  that  less  chlorine  is  formed  and  more  of  the  hydrochloric 
acid  passes  through  the  heated  apparatus  unchanged  (p.  110). 
By  the  use  of  a  catalyzer,  however,  the  reaction  at  lower  tem- 
peratures can  be  hastened  so  far  as  to  become  technically  pos- 
sible. 

Several  different  catalyzers  may  be  used,  but  Deacon  found 
that  copper  chloride,  in  the  form  obtained  by  saturating  pumice 
with  a  solution  of  the  salt  and  drying,  is  most  suitable.  With 
this  catalyzer  the  reaction  is  sufficiently  rapid  so  that,  at  345°, 
80  per  cent  of  the  hydrochloric  acid  can  be  oxidized  to  chlorine. 
The  process  does  not,  however,  compete  successfully  with  other 
processes  for  the  manufacture  of  chlorine. 

*  4.  The  Weldon  Process  for  Chlorine.  The  oxygen  of  the 
air  is  also  used,  indirectly,  as  the  oxidizing  agent  for  the  hydro- 
chloric acid  in  the  Weldon  process.  When  calcium  hydroxide 
(in  the  form  of  milk  of  lime)  is  added  to  the  solution  of  manganese 
chloride  obtained  in  the  preparation  of  chlorine  by  means  of 
manganese  dioxide,  a  precipitate  of  manganese  hydroxide  is 
formed : 

/Cl  /O— H  /OH  /Cl 

Mn<      +Ca<  =Mn<        +  Ca< 

XC1  XO— H  XOH  XC1 

or 

MnCl2  +  Ca(OH)2  =  Mn(OH)2  +  CaCl2 

The  manganese  hydroxide,  if  exposed  to  the  air,  takes  up 

OH 
oxygen  and  forms  hydrated  manganese  dioxide,  O  =  Mn( 

XOH 

(or  MnO2.H2O).  The  oxidation  is  hastened,  practically,  by 
passing  air  and  steam  into  the  mixture.  The  addition  of  the 
oxygen  to  the  manganese  causes  the  hydrogen  of  the  hydroxyl 
groups  to  become  acid  in  character  (see  p.  206),  and  the  com- 


104  A  TEXTBOOK  OF  CHEMISTRY 

pound  reacts  with  more  of  the  calcium  hydroxide  to  form  cal- 
cium manganite,  CaMnOa : 

/O— H  /OH  /Ov 

O=Mr<  +  Ca<         =  O=Mr<      >Ca  +  2  H2O 

X0— H          \)H  XX 

Calcium 
Manganite 

The  calcium  manganite  is  insoluble  and  may  be  easily  sepa- 
rated from  the  solution  of  calcium  chloride.  On  treatment  with 
hydrochloric  acid  it  acts  in  the  same  way  as  a  mixture  of  man- 
ganese dioxide  and  lime  would. 

CaMnO3  +  6  HC1  =  CaCl2  +  MnCl2  +  C12  +  3  H2O 

The  solution  of  manganese  chloride  may,  of  course,  be  treated 
with  milk  of  lime,  air  and  steam  and  the  cycle  repeated  indefi- 
nitely. This  process  was  used  for  many  years  in  manufacturing 
chlorine  and  bleaching  powder  on  a  large  scale,  but  it  has  now 
been  replaced  almost  entirely  by  the  electrolytic  processes, 
which  are  simpler  and  more  direct. 

Properties  of  Chlorine.  Chlorine  is  a  greenish  yellow  gas 
about  two  and  one  half  times  (££,  see  p.  95)  as  heavy  as  air.  It 
has  a  characteristic  odor  and,  even  when  diluted  with  a  large 
volume  of  air,  attacks  the  nose  and  lungs  strongly,  producing  the 
effect  of  a  severe  cold.  The  best  antidote  is  to  breathe,  at  once, 
the  vapors  of  strong  alcohol.  A  larger  quantity  of  the  gas  acts 
as  a  violent  poison  and  may  produce  fatal  effects. 

Water  dissolves  about  twice  its  volume  of  the  gas  and  for 
laboratory  experiments  it  is  usually  collected  by  displacement 
of  air,  in  upright  jars.  It  is  less  soluble  in  a  concentrated  solu- 
tion of  salt. 

Chlorine  may  be  condensed  to  a  liquid  by  cold  or  by  pressure. 
The  liquid  boils  at  —  33.6°  under  atmospheric  pressure  and 
freezes  at  —  102°.  The  vapor  pressure  at  0°  is  3.66  atmospheres. 

Chlorine  is  a  very  active  element  and  forms  compounds  with 
all  of  the  elements  except  fluorine  and  those  of  the  argon  family. 
With  many  of  the  elements  it  will  combine  directly  and  rapidly 


CHLORINE  105 

at  ordinary  temperatures  and  with  others  it  combines  at  a  much 
lower  temperature  than  does  oxygen. 

Chlorine  and  hydrogen  combine  too  slowly  for  the  rate  to  be 
measured,  if  the  mixture  of  the  gases  is  kept  in  the  dark,  but  the 
mixture  will  explode  at  a  much  lower  temperature  than  mix- 
tures of  oxygen  and  hydrogen.  If  the  mixture  is  exposed  to 
diffused  daylight,  the  elements  combine  slowly  at  ordinary 
temperatures,  and  if  exposed  to  bright  sunlight  or  to  the  light 
from  burning  magnesium,  the  combination  is  so  rapid  as  to  pro- 
duce a  violent  explosion.  The  effect  is  produced  by  light  and 
not  by  heat,  and  the  rays  of  light  at  the  violet  end  of  the  spec- 
trum are  especially  effective,  just  as  the  same  light  rays  are  most 
effective  in  producing  changes  in  photographic  plates.  Evi- 
dently the  light  vibrations  set  up  or  increase  some  kind  of  vibra- 
tion within  the  molecules  or  atoms  of  chlorine,  which  makes 
these  more  active,  but  we  can  form  only  a  very  vague  idea  of 
the  mechanism  of  the  action. 

Chlorine  containing  a  minute  amount  of  moisture  will  attack 
almost  all  of  the  metals,  even  at  ordinary  temperatures,  and 
more  rapidly  on  gentle  warming.  A  strip  of  copper,  if  warmed 
and  held  in  the  gas,  takes  fire  and  burns  to  cuprous  chloride, 
Cu2Cl2,  which  melts  and  runs  from  the  end  of  the  piece.  Dutch 
metal,  or  false  gold  leaf,  burns  with  a  flash.  It  contains  copper 
and  zinc  and  gives  cuprous  chloride,  Cu2Cl2,  and  zinc  chloride, 
ZnCl2-  Powdered  antimony  burns  with  brilliant  flashes  and 
gives  antimony  pentachloride,  SbCls.  Phosphorus  takes  fire  in 
chlorine,  burning  to  the  liquid  phosphorus  trichloride,  PCla,  if 
the  phosphorus  is  in  excess,  or  to  the  solid  phosphorus  penta- 
chloride, PCls,  if  the  chlorine  is  in  excess. 

If  chlorine  gas  is  carefully-  dried  with  phosphorus  pentoxide, 
it  will  not  act  on  copper,  iron  or  other  metals.  This  may  be 
shown  by  passing  the  dry  chlorine  into  a  flask  containing  dry 
Dutch  metal.  The  metal  will  remain  perfectly  bright,  but  the 
introduction  of  the  slightest  trace  of  moisture  will  cause  the  im- 
mediate combination  of  the  chlorine  with  the  leaf.  The  reason 
for  this  catalytic  effect  of  the  water  is  not  understood.  Because 


106  A  TEXTBOOK  OF  CHEMISTRY 

of  this  property,  dry,  liquid  chlorine  is  kept  and  sold  in  strong 
steel  cylinders. 

If  a  piece  of  tissue  paper,  which  has  been  dipped  in  warm  tur- 
pentine, CioHie,  is  thrust  into  a  jar  of  chlorine,  the  turpentine 
will  usually  take  fire  and  burn  with  a  very  smoky,  red  flame, 
giving  hydrochloric  acid  and  carbon  : 

Ci0H16  +  8  C12  =  10  C  +  16  HC1 

Chlorine  and  Water.  Bleaching.  A  solution  of  chlorine  in 
water  has  the  same  greenish  yellow  color  which  is  characteristic 
of  the  gas.  Such  a  solution  apparently  contains  most  of  the 
chlorine  as  such,  but  it  has  been  shown  (Jakowin,  Z.  physik.  Chem. 
29,  613)  that  a  small  amount  of  the  chlorine  reacts  with  the 
water : 

Cl— Cl  +  H— O— H  ^±  HC1  +  Cl— O— H 

Hypochlorous 
Acid 

The  reaction  is  reversible,  with  the  equilibrium  far  toward 
the  side  giving  chlorine  and  water,  and  very  little  hypochlorous 
acid  is  present  in  the  solution.  If  the  solution  is  exposed  to 
light,  however,  the  hypochlorous  acid  decomposes  in  two  ways, 
giving  either  chloric  and  hydrochloric  acids  or  oxygen  and  hy- 
drochloric acid : 

2  HC1O  +  HC1O  =  2  HC1  +  HC1O3 

Chloric 
Acid 

2  HC1O  =  2  HC1  +  O2 

It  is  very  noticeable  that  the  compounds  of  chlorine  with  oxy- 
gen become  more  stable  as  the  amount  of  oxygen  in  them  in- 
creases. As  the  two  reactions  progress,  the  color  of  the  chlorine 
gradually  disappears. 

If  dry  litmus  paper  or  a  dry  piece  of  colored  calico  is  placed 
in  dry  chlorine,  the  color  is  affected  very  slowly,  if  at  all,  but  if 
the  paper  or  cloth  is  moistened,  the  color  will  be  bleached  very 
quickly.  The  effect  of  the  water  may  be  in  part  similar  to  the 
action  of  moisture  in  causing  chlorine  to  combine  with  metals, 


PHASES  107 

but  the  chlorine  also  reacts  with  the  water  giving  hypochlorous 
acid,  HC1O,  and  this  oxidizes  and  destroys  the  coloring  matter. 
*  Chlorine  Hydrate.  Phases.  If  chlorine  is  passed  into  water 
which  is  cooled  to  0°,  a  crystalline  compound,  chlorine  hydrate, 
C12.8  H2O,  separates.  If  this  is  warmed  under  atmospheric 
pressure,  it  decomposes  at  9.6° ;  but  if  the  pressure  of  the  chlo- 
rine is  increased,  it  may  exist  at  higher  temperatures,  or  if  the 
pressure  is  lessened,  it  will  decompose  at  a  lower  temperature. 
We  have,  in  this  case,  a  system  of  three  phases,  solid,  liquid  and 
gas,  which  can  exist  over  a  range  of  several  degrees  of  tempera- 
ture. This  is  true  of  any  other  system  containing  two  com- 
ponents. Such  a  system  with  two  phases  may  still  have  two 
degrees  of  freedom  (see  p.  77)  —  freedom  to  change  in  tempera- 
ture and  freedom  to  change  in  pressure.  The  addition  of  a 
second  component,  chlorine,  increases  the  number  of  degrees 
of  freedom  for  a  given  number  of  phases.  If  the  system  is  cooled 
to  —  0.24,  ice  will  separate  from  the  solution  of  chlorine,  as  well 
as  chlorine  hydrate.  When  this  occurs,  there  will  be  four  phases 
present,  liquid,  ice,  chlorine  hydrate  and  vapor  or  gas.  The  pres- 
sure will  also  be  fixed  at  244  mm.  No  change  in  either  tempera- 
ture or  pressure  can  occur  without  the  disappearance  of  one  of 
the  phases.  With  two  components  and  four  phases,  there  is  no 
freedom.  A  further  study  of  cases  in  which  there  are  two  or 
more  components  leads  to  the  conclusion  that  the  number  of 
phases  and  number  of  degrees  of  freedom  together  are  equal  to 
the  number  of  components  increased  by  two,  or : 

p-f  p  =  C  +  2 

P  =  number  of  phases 
F  =  degrees  of  freedom 
C  =  number  of  components 

This  is  the  celebrated  "Phase  rule,"  which  was  discovered  by 
Willard  Gibbs  of  Yale  University.  It  applies  equally  well  to 
the  formation  and  decomposition  of  compounds,  as  of  chlorine 
hydrate  above,  and  to  changes  of  state,  as  from  ice  to  water  and 
vapor.  It  is  applicable  only  when  the  changes  in  state  are 


108  A  TEXTBOOK  OF  CHEMISTRY 

reversible  and  is  important  only  when  the  equilibrium  between 

the  different  phases  is  reached  within  a  measurable  time. 

Faraday  first  prepared  liquid  chlorine 
by  warming  chlorine  hydrate  in  a  bent 
tube  of  the  form  shown  in  Fig.  38. 
By  immersing  the  closed,  empty  end, 
B,  in  a  freezing  mixture  while  the 
chlorine  hydrate  in  the  end,  A,  was 
Fig.  38  warmed,  the  gas  liberated  by  the  de- 

composition   of    the    hydrate  exerted 

enough  pressure  to   cause   a    part    to    liquefy    in    the  cooled 

end. 

The  Heat  of  Combination  of  Chlorine  and  of  Oxygen  with 

Other  Elements.     The  following  are  the  heats  of  combination 

of  several  elements  with   equivalent  amounts  of  oxygen  and 

chlorine : 

H2  +     O     =  H2O  ( vapor)  +    58,000  calories 

H2  +     C12  =  2  HC1  +    44,000  calories 

2  Na  +     O     =  Na2O  +  100,000  calories 

2  Na  +     C12  =  2  NaCl  +  195,000  calories 

Zn  +     O     -  ZnO  +    85,300  calories 

Zn  +     C12  =  ZnCl2  +    97,200  calories 

Cu  +     O     =  CuO  +    37,200  calories 

Cu  +     C12  =  CuCl2  +    51,600  calories 

P2  -f  5  O     =  P2O5  +  370,000  calories 

.  P2  +  5  C12  =  2  PC15  +  210,000  calories 

The  heat  of  combination  with  chlorine  is  sometimes  greater, 
sometimes  less,  than  the  heat  of  combination  with  oxygen.  In 
general,  the  heat  of  combination  of  chlorine  seems  to  be  greater 
than  that  of  oxygen  in  combining  with  metals  and  less  than  that 
of  oxygen  in  combining  with  nonmetals. 

Equilibrium  in  Chemical  Reactions.  If  a  mixture  of  four 
volumes  of  hydrochloric  acid  with  one  volume  of  oxygen  is  passed 
slowly  through  a  tube  containing  cuprous  chloride  at  345  , 


EQUILIBRIUM 


109 


four  fifths  of  the  hydrochloric  acid  will  be  oxidized,  giving  chlo- 
rine and  water  in  accordance  with  the  reversible  reaction : 

4  HC1  +  02  =  2  C12  +  2  H20 

If  a  mixture  of  equal  volumes  of  chlorine  and  steam  is  passed 
slowly  through  the  tube  at  the  same  temperature,  one  fifth  of 
the  chlorine  will  be  converted  into  hydrochloric  acid.  In  other 
words  a  mixture  containing 

4  volumes  or  4  molecules  of  HC1 
1  volume   or  1  molecule   of  O2 
8  volumes  or  8  molecules  of  C12 
8  volumes  or  8  molecules  of  H2O 

will  be  in  equilibrium  and  will  not  change  its  composition  when 
heated  for  a  long  time  at  345°.  We  do  not  suppose  that  chemical 
action  ceases,  but  rather  that,  in  a  given  time,  just  as  many  atoms 


20  Vols.  HC1 1 
5  Vols.  O2     I 


[  4  Vols.  HC1 
1  Vol.  O2 
8  Vols.  Cl, 

( 8  Vols.  H2O 


Fig.  39 

of  chlorine  unite  with  hydrogen  to  form  hydrochloric  acid  as 
there  are  atoms  of  chlorine  separated  from  molecules  of  hydro- 
chloric acid.  In  this  way  the  total  number  of  molecules  of  each 


10  Vols.  C12 
10  Vols.  H2O 


[  4  Vols.  HC1 
J  1  Vol.  O2 
8  Vols.  C12 
8  Vols.  H20 


Fig.  40 


110  A  TEXTBOOK  OF  CHEMISTRY 

of  the  four  substances  will  remain  unchanged  after  equilibrium 
is  reached,  but  any  given  atom  may  frequently  change  its  state 
of  combination.  In  s"uch  a  case  we  may  think  of  two  opposing 
forces,  one  of  which  drives  the  reaction  to  the  right  and  the  other 
drives  it  to  the  left,  and  that  these  forces  are  in  equilibrium. 
The  force  driving  the  reaction  toward  the  right,  when  we  start 
with  hydrochloric  acid  and  oxygen,  must  be  much  stronger  than 
the  force  driving  the  reaction  toward  the  left,  when  we  start  with 
chlorine  and  steam.  This  is  probably  due,  in  part,  to  the  greater 
affinity  of  oxygen  for  hydrogen,  as  indicated  by  the  heat  of  com- 
bination given  in  the  last  paragraph,  but  it  is  also  connected  with 
the  change  in  volume  which  occurs  in  the  reaction  and  with  other 
factors  which  are  less  clearly  understood. 

The  reaction  proceeds  toward  the  right  with  the  evolution  of 
heat : 

H2     +  O     =     H2O  -f  58,000  calories 
H2     +  C12  =  2  HC1  +  44,000  calories 
Hence  2  HC1  +  O    =      H2O  +  C12  +  14,000  calories 

since  the  sum  of  the  reactions : 

H2     +  Cla  =  2  HC1 
and  2  HC1  +  O     =  H2O  +  C12 

must  give  the  same  amount  of  heat  as  the  reaction  H2  +  O  =  H2O, 
because  the  chlorine  is  in  the  same  condition  at  the  end  as  at 
the  beginning. 

Whenever  a  reversible  reaction  proceeds  with  evolution  of 
heat,  a  higher  temperature  always  shifts  the  equilibrium  in  the 
direction  to  cause  a  smaller  evolution  of  heat.  In  other  words, 
the  application  of  heat  always  helps  the  side  of  a  reversible  reac- 
tion in  which  heat  is  absorbed  and  retards  that  side  of  a  reaction 
in  which  heat  is  given  out. 

In  accordance  with  this  we  find  that  the  mixture  in  equilibrium 
at  384°  contains : 


EQUILIBRIUM  111 

4  volumes  or  4  molecules  of  HC1 
1  volume   or  1  molecule   of  C>2 
6  volumes  or  6  molecules  of  C12 
6  volumes  or  6  molecules  of  H^O 

This  means  that  while  four  fifths  of  the  hydrochloric  acid  can 
be  oxidized  to  chlorine  by  the  Deacon  process  at  345°,  only  Ijiree 
fourths  of  it  can  be  oxidized  at  384°.  It  is  this  very  unfavor- 
able effect  of  an  increase  in  the  temperature  which  makes  it 
necessary  to  use  a  catalyzer  and  work  at  as  low  a  temperature 
as  possible. 

*  Principle  of  van't  Hoff-Le  Chatelier.  —  As  has  been  stated 
above,  an  increase  in  temperature  displaces  any  equilibrium  in 
the  direction  in  which  heat  is  absorbed  and  an  increase  in 
pressure  displaces  any  equilibrium  in  the  direction  in  which  the 
volume  decreases.  These  are  special  cases  of  the  principle  of 
van't  Hoff-Le  Chatelier,  which  is  that  every  force  applied  to  a 
system  which  is  in  equilibrium  causes  a  change  which  tends  to 
resist  the  force  that  is  applied.  Thus,  if  pressure  is  applied  to 
ice,  a  small  amount  of  the  ice  will  melt,  but  in  melting  it  will 
absorb  heat,  the  temperature  will  fall  and  this  will  tend  to 
stop  the  melting.  Or,  if  dry  air  is  blown  over  the  surface  of 
water,  it  will  cause  the  water  to  evaporate;  but  as  it  evap- 
orates, the  temperature  will  fall  and  the  lower  vapor  pressure 
will  tend  to  stop  the  evaporation. 

In  accordance  with  the  law,  if  hydrochloric  acid  and  oxygen 
are  brought  together  at  345°,  heat  will  be  evolved  as  they  react, 
and  this  will  tend  to  stop  the  reaction.  On  the  other  hand,  if 
steam  and  chlorine  are  brought  together  at  345°,  heat  will  be 
absorbed  as  they  react,  and  this  will  tend  to  increase  the  re- 
action. 

It  is  necessary  to  distinguish  between  the  effect  of  an  in- 
crease in  temperature  to  increase  the  speed  of  a  reaction, 
which  seems  to  be  universal,  and  the  tendency  to  cause  a 
reversal  of  those  reactions  in  which  heat  is  evolved.  The  speed 
of  the  combination  of  oxygen  and  hydrogen  increases  rapidly 
with  the  temperature  and  becomes  explosive  at  a  very  mod- 


112  A  TEXTBOOK  OF  CHEMISTRY 

erate  heat.  The  reverse  reaction  by  which  water  dissociates 
into  oxygen  and  hydrogen  increases,  however,  with  the  tem- 
perature in  accordance  with  the  theorem  of  van't  Hoff-Le 
Chatelier. 

If  we  could  understand  fully  the  mechanism  of  all  physical 
and  chemical  processes,  it  seems  likely  that  we  should  find  that 
this*  principle  has  its  foundation  in  Newton's  law  that  for 
every  action  there  is  an  equal  and  opposite  reaction. 

Effect  of  Water  on  Chlorides.  lonization.  When  hydro- 
chloric acid  or  such  chlorides  of  the  metals  as  sodium  chloride, 
NaCl,  zinc  chloride,  ZnCl2,  or  copper  chloride,  CuCl2,  dissolve 
in  water,  several  different  lines  of  evidence  indicate  that  these 
compounds  separate  more  or  less  completely  into  chloride  ions, 
Cl~~,  bearing  a  negative  charge  of  electricity  and  hydrogen, 
H+,  or  metallic  ions,  Na+,  Zn++,  or  Cu++,  bearing  a  positive 
charge,  or,  if  the  atom  is  bivalent  or  trivalent,  positive  charges 
of  electricity.  The  evidence  for  this  view  is,  in  part,  as  fol- 
lows :  , 

1.  It  is  found  that  if  substances  which  are  not  electrolytes  are 
dissolved  in  water,  the  freezing  point  is  lowered  in  direct  propor- 
tion to  the  amount  dissolved  and  in  inverse  proportion  to  the 
molecular  weight  of  the  solute.  An  aqueous  solution  of  alcohol 
(C2H6O,  molecular  weight,  46)  or  of  sugar  (C^B^On,  molec- 
ular weight,  342)  is  almost  as  poor  a  conductor  as  pure  water. 
A  solution  containing  46  milligrams  of  alcohol  dissolved  in  10  cc. 
of  water  will  freeze  at  —  0.184°.  A  solution  of  sugar  containing 
342  milligrams  in  10  cc.  of  water  will  freeze  at  —  0.188.°  One  of 
these  solutions  must  contain  the  same  number  of  molecules  as 
the  other,  and  it  is  evident  that  the  lowering  of  the  freezing  point 
is  proportional  to  the  number  of  molecules  of  the  solute  in  a  given 
volume  of  the  solvent.  But  if  we  dissolve  58.5  milligrams  of 
salt  (NaCl,  molecular  weight,  58.5)  in  10  cc.  of  water,  the  solution 
will  freeze  at  —  0.349°.  According  to  the  law  just  stated,  the 
freezing  point  indicates  nearly  twice  as  many  molecules  as  there 
should  be.  The  simplest  explanation  of  this  fact  is  that  the 
sodium  chloride  separates  into  sodium  (Na+)  and  chloride  (Cl~) 


IONIZATION 


113 


ions  in  the  solution  and  that  these,  so  far  as  this  law  is  con- 
cerned, act  as  independent  molecules. 

In  a  similar  manner,  a  solution  containing  36.5  milligrams  of 
hydrochloric  acid  (HC1,  molecular  weight,  36.5)  in  10  cc.  of  water 
freezes  at  —  0.355°,  indicating  that  it  is  largely  separated  into 
hydrogen  (H~)  and  chloride  (Cl~)  ions. 

2.  If  charged  strips  of  metal,  as,  for  instance,  pieces  of  platinum 
connected  with  the  poles  of  an  electric  battery,  are  dipped  in  a 
solution  of  hydrochloric  acid,  the  chlorine  atoms  are  attracted 
by  the  positive  electrode  (anode)  and  the  hydrogen  atoms  are 
attracted  by  the  negative  electrode  (cathode)  and  there  is  a 
motion  of  the  ions  throughout  the  solution.  This  is  called  the 
migration  of  the  ions.  The  rate  of  migration  varies  for  different 
ions  and  with  the  same  current  it  is  quite  different,  the  hydrogen 
ions  moving  nearly  five  times  as  fast  as  the  chloride  ions. 
This  effect  can  be  shown  by  passing  an  electrical  current 
from  a  silver  anode  to  a  platinum  cathode  through  hydro- 
chloric acid  in  the  U-tube  (Fig.  41).  A  silver  anode  is  used 
because  it  will  combine  quantitatively  with  the  chlorine  which 
is  liberated,  and  remove  it  from  the  solution. 


The  solution  on  this  side 

contains :  Cathode 

At  first,  j^\- 

73  mg.  HC1  = 
2  milligram  atoms  of  H 
2  milligram  atoms  of  Cl 
At  end, 

66.5  mg.  HC1  = 
1.82    milligram    atoms 

of  H 
1.82  milligram  atoms  of 

Cl 
1  milligram  atom  of  H 

is  liberated 

Amount  of  hydrogen 
transferred  across  the 
line  CD  =  0.82  mg. 
atoms. 


Anode 


4 

V 

i 

g 

y 

y 

\ 

/ 

^^4-^ 

D 

Fig.  41 

The  solution  on  this  side 
t       contains : 
At  first, 
73  mg.  HC1  = 
2  milligram  atoms  of  H 
2  milligram  atoms  of  Cl 
At  end, 
43  mg.  HC1  •= 
1.18  milligram  atoms  of 

H 
1.18  milligram  atoms  of 

Cl 
1  milligram  atom  of  Cl 

combines  with  anode 
Amount  of  chlorine 
transferred  across  the 
line  CD  =  0.18  mg. 
atoms. 


114 


A  TEXTBOOK  OF  CHEMISTRY 


At  the  beginning  of  the  experiment  the  concentration  of  the 
hydrochloric  acid  is  the  same  in  both  arms  of  the  tube ;  but  after 
decomposing  a  part  of  the  hydrochloric  acid  by  passing  the  cur- 
rent for  some  time,  it  will  be  found  that  while  the  concentration 
of  the  acid  has  decreased  on  both  sides,  the  amount  of  acid  on 
the  cathode  side  has  become  much  greater  than  that  on  the  anode 
side.  Since  the  number  of  hydrogen  atoms  liberated  at  the 
cathode  must  be  exactly  the  same  as  the  number  of  chlorine 
atoms  which  combine  with  the  silver  anode,  the  greater  amount 
of  acid  on  the  cathode  side  must  be  due  to  the  fact  that  the 
hydrogen  ions  migrate  faster  toward  the  cathode  than  the 
chloride  ions  migrate  toward  the  anode.  This  will  be  clear  from 
an  examination  of  the  figure  and  the  accompanying  statement 


f 


Fig.  42 

about  the  composition  of  the  solution  at  the  beginning  and  end 
of  the  experiment. 

3.  If  a  solution  of  sodium  iodide  is  subjected  to  a  powerful 
centrifugal  force,  the  heavier  iodide  ions  may  be  separated  to  a 
slight  extent  from  the  sodium  ions  (Tolman,  J.  Am.  Chem. 
Soc.  33,  121).  Similar  experiments  were  tried  with  hydriodic 


IONIZATION  115 


acid,  lithium  iodide  and  potassium  iodide.  By  whirling  solu- 
tions of  these  substances  in  the  apparatus  shown  diagrammati- 
cally  in  Fig.  42  the  heavy  iodide  ions  were  thrown  toward  the 
outside,  giving  a  negative  electrical  charge  to  the  solution  at 
that  end  of  the  tube  and  a  positive  charge  at  the  inner  end. 

All  substances  which  are  electrolytes  are  supposed  to  separate 
more  or  less  into  ions  in  aqueous  solutions.  Thus  sodium  nitrate, 
NaNOs,  separates  into  sodium  ions  (Na+)  and  nitrate  ions 
(NOs~) ;  sulfuric  acid,  H2SO4,  may  separate  into  two  hydrogen 
ions  (H+,  H+),  and  the  sulfate  ion  (SO4  ),  or  it  may  separate, 
in  part,  only  into  a  single  hydrogen  ion  (H+)  and  the  acid  sul- 
fate ion  (HSO4+).  When  solutions  containing  electrolytes  are 
mixed,  the  reactions  which  occur  are,  in  most  cases,  a  simple 
exchange  of  ions,  and  the  groups  of  atoms  which  form  the  ions 
remain  unbroken.  Thus  in  the  reaction  : 

AgNO3  +  HC1  =  AgCl  +  HNO3 

the  nitrate  ion  (NOa~)  passes  from  one  compound  to  the  other 
without  any  change. 

Such  reactions  are  always  reversible,  and  the  equilibrium  is 
frequently  displaced  to  one  side  or  the  other  because  one  of  the 
compounds  is  volatile  or  difficultly  soluble.  Thus  it  has  been 
pointed  out  that  in  the  reaction : 

NaCl  +  H2S04  ^±  HC1  +  NaHSO4 

if  concentrated  sulfuric  acid  is  added  to  salt  the  equilibrium  will 
be  far  to  the  right  because  the  hydrochloric  acid  is  a  gas  and 
escapes ;  while  if  a  concentrated  solution  of  'hydrochloric  acid 
is  added  to  a  concentrated  solution  of  acid  sodium  sulfate,  the 
equilibrium  may  be  carried  to  the  left,  because  sodium  chloride 
precipitates. 

Effect  of  Water  on  Chlorides.  Hydrolysis.  When  water  is 
brought  in  contact  with  a  chloride  of  a  nonmetallic  element,  the 
effect  is  very  different.  The  chloride  reacts  with  the  water  as 
though  the  water  were  composed  of  two  parts,  hydrogen,  H,  and 
hydroxyl,  OH. 


116  A  TEXTBOOK  OF  CHEMISTRY 


C\        HOH  ,OR  OH 

C1  +     HOH  =  3  HC1  +  P^OH  or  O=P^OH 
C1         HOH  XOH  XH 

Phosphorous  Acid. 

01  ,OH 

/OH  O 

+  5  HOH  =  5  HC1  +  P^-OH  =  P^OH  +  H2O. 

yci  YOH       \OH 

\C1  \OH          \OH 

Phosphoric  Acid 

This  sort  of  double  decomposition  with  water  is  called  hy- 
drolysis. The  hydroxyl  compounds  which  are  formed  are  acids. 
In  most  reactions  between  these  acids  and  other  compounds, 
the  oxygen  is  held  by  the  nonmetallic  element,  while  the  hy- 
drogen may  be  easily  replaced  by  metals.  Thus  with  sodium 
hydroxide  we  have  the  reaction  : 

H3PO4  +  3  NaOH  =  Na3PO4  +  3  HOH 

The  division  between  the  metallic  and  nonmetallic  elements 
in  the  conduct  of  the  chlorides  is  not  a  sharp  one.  While  sodium 
or  potassium  chlorides  are  only  ionized  in  solution  and  on 
evaporation  of  the  water  the  ions  recombine  without  any  loss 
of  hydrochloric  acid,  and  phosphorus  pentachloride  is  completely 
decomposed  by  water  and  on  evaporation  of  the  solution  the  hy- 
drochloric acid  will  escape  entirely,  leaving  phosphoric  acid, 
there  are  many  other  chlorides,  like  ferric  chloride  and  aluminium 
chloride,  which  partly  ionize  and  partly  hydrolyze  in  solution. 

EXERCISES 

1.  If  a  mixture  of  salt  and  manganese  dioxide  is  treated  with  sulfuric 
acid,  the  products  will  be  manganous  sulfate,  sodium  sulfate,  water 
and  chlorine.     Write  the  equation. 

2.  How  much  hydrochloric  acid  will  be  required  to  give  10  liters  of 
chlorine  by  the  Deacon  process,  assuming  that  80  per  cent  is  oxidized  to 
chlorine  ?     How  much  hydrochloric  acid  will  be  required  by  the  second 
stage  of  the  Weldon  process  ? 


CHLORINE  117 

3.  What  per  cent  of  the  chlorine  in  hydrochloric  acid  is  liberated 
when  the  acid  acts  on  manganese  dioxide  ?     What  per  cent  when  it  acts 
on  potassium  permanganate?     What  .per  cent  when  it  acts  on  calcium 
manganite  ? 

4.  What  is  the  weight  of  chlorine  absorbed  by  one  liter  of  water  ? 

5.  Assuming  that  air  contains  21  per  cent  of  oxygen  (by  volume), 
how  many  volumes  of  hydrochloric  acid  should  be  mixed  with  100  vol- 
umes of  air  for  the  Deacon  process  ?     What  per  cent  of  free  chlorine  will 
the  resulting  mixture  contain,  after  the  reaction,  if  there  is  an  oxidiza- 
tion of  80  per  cent  ? 


CHAPTER  VIII 

HYDROCHLORIC   ACID.     OXIDES   AND   OXYACIDS   OF 
CHLORINE 

Hydrochloric  Acid.  The  explosive  combustion  of  hydrogen 
and  chlorine  under  the  influence  of  light  has  been  mentioned. 

Hydrogen  may  be  burned  in  a  jar 

of  chlorine  or  chlorine  may  be 

burned  in  hydrogen,  Figs.  43  and 

44.     In   each   case  hydrochloric 

acid  is  formed.     It  may  be  pre- 
pared more  easily  by  pouring  a 

mixture  of  9  parts  (by  weight)  of 

concentrated  sulfuric  acid  with  2 

parts  of  water  on  common  salt, 
Fig.  43         NaCL     Concentrated  acid  might 

be  used,  but  the  mixture  with 
salt  froths  badly,  while  the  slightly  diluted  acid  does  not  froth : 

NaCl      +   H2SO4  :£     NaHSO4      +   HC1 

Sodium  Acid  Sodium 

Chloride  Sulfate 

The  compound  NaHSO4  is  called  acid  sodium  sulfate  because 
it  still  contains  an  acid  hydrogen  atom  which  can  be  replaced 
by  a  metal.  Thus  if  more  salt  is  added  and  the  mixture  is 
warmed,  the  reaction : 

NaHSO4  +  NaCl  ^±  Na2SO4  +  HC1 

will  occur. 

These  reactions  are  reversible  and  would  be  very  far  from  com- 
plete in  either  direction  if  all  of  the  substances  remained  mixed 
together.  But  as  soon  as  the  salt  and  sulfuric  acid  are  mixed, 
hydrochloric  acid  begins  to  escape  as  a  gas.  When  this  occurs, 

118 


HYDROCHLORIC  ACID  119 

the  acid  which  has  gone  can  no  longer  have  any  effect  in  driving 
the  reaction  in  the  oppo'site  direction,  and  a  new  quantity  of  the 
sulfuric  acid  will  act  on  the  salt.  In  this  way,  if  the  mixture  is 
warmed,  the  reaction  may  finally  be  made  practically  complete. 
The  equilibrium  is  displaced  in  the  direction  toward  the  forma- 
tion of  the  product  which  is  continually  removed  from  the 
mixture. 

That  the  reaction  is  reversible  may  be  easily  shown  by  adding 
a  concentrated  solution  of  hydrochloric  acid  to  a  strong  solu- 
tion of  acid  sodium  sulfate.  A  copious  precipitate  of  sodium 
chloride  will  be  formed  : 

HC1  +  NaH3O4  ^±  NaCl  +  H2SO4 

In  this  case,  as  the  sodium  chloride  precipitates  it  can  no  longer 
act  on  the  substances  remaining  in  solution,  and  the  equilibrium 
is  displaced  toward  the  formation  of  the  compound  which  is 
precipitated.  Not  many  years  ago  it  was  quite  common  to  say 
that  sulfuric  acid  is  stronger  than  hydrochloric  acid  and  so  expels 
hydrochloric  acid  from  its  salts.  The  experiments  described 
show  that  either  acid  may  expel  the  other,  and  that  the  direction 
in  which  the  reaction  goes  depends  on  the  volatility  or  insolu- 
bility of  the  compounds  formed  and  on  the  relative  amounts  of 
the  reacting  substances  as  well  as  upon  the  relative  affinities 
of  the  chlorine  and  of  the  sulfate  radical  for  the  metal,  and  that 
the  first  three  factors  are  frequently  more  important  than  the 
last. 

Properties  of  Hydrochloric  Acid.  Hydrochloric  acid  is  a 
colorless  gas  which  may  be  condensed  to  a  liquid  by  cold  and 
pressure.  The  liquid  boils  at  -  83.7°  and  freezes  at  -  110°. 
What  is  the  weight  of  22.4  liters  of  the  gas  ?  What  is  the  density 
as  compared  with  air  ? 

Water  at  0°  will  absorb  503  volumes  of  the  gas.  If  the  solu- 
tion is  boiled,  more  hydrochloric  acid  than  water  escapes  at  first, 
and  the  temperature  gradually  rises  till  a  boiling  point  of  110° 
is  reached.  After  that,  the  portion  which  distils  over  and  that 
which  remains  behind  will  have  the  same  composition,  contain- 


120 


A  TEXTBOOK  OF  CHEMISTRY 


ing  20.2  per  cent  of  the  acid.  If  an  acid  which  contains  less 
than  20.2  per  cent  is  boiled,  the  boiling  point  will  be  below  110° 
and  the  portion  distilling  over  will  contain  less  acid  than  that 
which  remains.  If  the  distillation  is  continued,  the  temperature 
will  gradually  rise  to  110°,  and  after  that  a  mixture  of  constant 
composition  (20.2  per  cent)  will  distill  as  before. 

The  ratio  between  the  volumes  of  hydrochloric  acid  and  of 
the  hydrogen  which  it  contains  may  be  demonstrated  roughly 
by  filling  a  dry  tube  with  the  gas,  pouring  in  a  few  cubic  centi- 
meters of  liquid  sodium  amalgam,  inserting  a  rubber  stopper 
quickly  and  shaking  vigorously.  On  opening  the  tube  with  the 
mouth  below  the  surface  of  water  in  a  beaker  the  water  will  rise 
and  fill  the  tube  one  half  full.  Does 
the  experiment  demonstrate  that  hy- 
drochloric acid  is  composed  of  equal 
volumes  of  hydrogen  and  chlorine? 
What  would  be  the  result  if  a  similar 
experiment  could  be  tried  with  steam 
and  all  of  the  hydrogen  of  the  steam 
were  replaced  by  the  metal  ? 

The  composition  of  hydrochloric  acid 
by  volume  may  be  demonstrated  by 
the  electrolysis  of  a  strong  solution  of 
the  acid,  using  carbon  electrodes.     The 
volumes  of  hydrogen  and  of  chlorine 
liberated  at  the  two  electrodes  will  be 
nearly  equal,  if  a  moderately  strong  cur- 
rent is  used  and  the  current  is  continued 
till  the  solution  around  the  anode  is 
saturated  with  chlorine.    A  suitable  ap- 
paratus is  shown  in  Fig.  45 .   See  Brown- 
lee,  J.  Am.  Chem.  Soc.  29,  237. 
For  most  laboratory  purposes  to  which  hydrochloric  acid  is 
applied  the  solution  in  water  is  used.     The  most  important 
chemical  properties  are  :  • 

1.  Reaction  with  Metals.     With  many  metals  the  hydrogen 


Fig.  45 


HYDROCHLORIC  ACID  121 


is  displaced  by  the  metal  and  chlorides  are  formed,  which  dis- 
solve in  the  water.  Thus,  sodium,  magnesium,  zinc,  iron, 
aluminium  and  tin  give  sodium  chloride,  NaCl,  magnesium  chlo- 
ride, MgCl2,  zinc  chloride,  ZnCl2,  ferrous  chloride,  FeCl2,  alumin- 
ium chloride,  A1C13,  and  stannous  chloride,  SnCl2.  What  are 
the  reactions  for  the  formation  of  these  chlorides  ?  It  is  worthy 
of  notice  that  metals  like  iron  and  tin,  which  form  two  chlorides, 
give  the  lower  chloride  when  the  metals  are  dissolved  in  a  solu- 
tion of  hydrochloric  acid. 

2.  Reaction  with  Hydroxides  of  Metals.  Hydrochloric  acid 
reacts  with  hydroxides  of  the  metals,  forming  chlorides  and 
water : 

HC1  +    NaOH       =     NaCl  +      HOH 

2  HC1  +  Fe(OH)2     =     FeCl2  +  2  HOH 

Ferrous  Ferrous 

Hydroxide  Chloride 

3  HC1  +  Fe(OH)3     =     FeCl3  +  3  HOH 

Ferric  Ferric 

Hydroxide  Chloride 

4  HC1  +  Sn(OH)4     =     SnCl4  +  4  HOH 

Stannic  Stannic 

Hydroxide  Chloride 

In  these  reactions  the  separation  of  the  metallic  hydroxide  is 
between  the  metal  and  hydroxyl,  just  as  the  separation  of  the 
acid  is  between  the  hydrogen  and  chlorine.  Compounds  which 
react  in  this  manner  are  called  bases,  the  presence  of  a  hydroxyl 
group,  OH,  which  separates  easily,  being  characteristic  of  a  base, 
as  the  presence  of  hydrogen  which  separates  easily  is  character- 
istic of  an  acid.  Since  hydrogen  and  hydro*xyl  have  a  strong 
affinity  for  each  other  and  separate  only  to  a  trifling  extent  in 
solutions  or  in  pure  water,  bases  and  acids  neutralize  each  other 
by  the  union  of  the  hydrogen  of  the  acid  with  the  hydroxyl  of  the 
base.  The  compound  formed  by  the  union  of  the  metal  with  the 
chlorine  or  with  the  acid  radical  is  called  a  salt.  In  each  case, 
for  the  formation  of  a  normal  salt  there  must  be  as  many  hy- 
droxyl groups  in  the  base  as  there  are  hydrogen  atoms  in  the 
acick 


122  A  TEXTBOOK  OF  CHEMISTRY 

3.  Reaction  with  Oxides  of  Metals.     Some  oxides  of  metals 
also  react  with  hydrochloric  acid  to  form  salts  and  water : 

ZnO  +  2  HC1  =  ZnCl2  +  H2O 

4.  Reaction  with  Oxidizing  Agents.     With  oxidizing  agents 
hydrochloric  acid  is  oxidized  to  water  and  chlorine,  the  oxidizing 
agent  being  at  the   same   time   reduced.     In   such   reactions 
chlorides  which  contain  two  atoms  of  chlorine  are  in  the  same 
degree  of  oxidation  as  those  which  contain  one  atom  of  oxygen, 
since  two  atoms  of  chlorine  replace  one  atom  of  oxygen  in  com- 
bination with  hydrogen.     Thus  manganous  oxide,  MnO,  is  in 
the  same  state  of  oxidation  as  manganous  chloride,  MnCl2, 
and  ferric  chloride,  FeCl3,  corresponds  in  oxidation  to  either 
ferric  oxide,  Fe2O3,  or  ferric  hydroxide,  Fe(OH)3,  while   man- 
ganese dioxide,  MnO2,  is  in  a  higher  state  of  oxidation  than  man- 
ganese chloride,  MnCl2. 

The  reactions  between  hydrochloric  acid  and  manganese  diox- 
ide, MnO2,  potassium  permanganate,  KMnC>4,  and  calcium  man- 
ganite,  CaMnO3,  have  been  given.  Similar  reactions  take  place 
with  lead  dioxide,  PbO2,  and  red  lead,  Pb3O4,  which  are  re- 
duced to  lead  chloride,  PbCl2 ;  also  with  potassium  dichromate, 
K2Cr2O7,  the  chromium  being  reduced  to  chromic  chloride, 
CrCl3,  while  the  potassium,  which  is  univalent,  does  not  change 
its  state  of  oxidation.  These  equations  should  be  written  by  the 
student  as  an  aid  to  an  understanding  of  reactions  of  this  type 
and  also  to  give  practice  in  writing  equations  correctly  by  devel- 
oping them  from  a  knowledge  of  the  compounds  formed  instead 
of  as  a  matter  of  memory. 

Why  does  not  barium  peroxide,  BaO2,  give  chlorine  when 
treated  with  hydrochloric  acid  ? 

Indicators.  A  number  of  organic  compounds  are  known 
which  have  one  color  in  an  acid  solution,  that  is  in  a  solution 
containing  hydrogen  ions,  H+,  and  another  color  in  an  alkaline 
solution,  that  is,  in  a  solution  containing  hydroxide  ions,  OH~. 
More  strictly  speaking,  such  compounds  are,  in  reality,  each  of 
them,  two  different  compounds  so  related  that  hydrogen  ions 


OXIDES  AND   OXYACIDS  OF  CHLORINE 


123 


will  change  the  first  into  the  second,  and  hydroxide  ions  will 
change  the  second  into  the  first.  Thus  litmus  is  a  red  compound 
in  an  acid  solution,  and  a  blue  compound  in  an  alkaline  solution. 
Some  of  the  common  indicators  are : 


NAME 

COLOR  IN 
ACID  SOLUTIONS 

COLOR  IN 

ALKALINE  SOLUTIONS 

Litmus 

Red 

Blue 

Phenolphthalei  n 

Colorless 

Red 

Methyl  orange 

Rose  red 

Yellow 

Methyl  red 

Red 

Yellow 

Congo  red 

Red 

Blue 

Oxides  and  Oxygen  Acids  of  Chlorine.  Nomenclature 
Chlorine  forms  three  oxides  and  four  acids  containing  oxygen : 

„,  ~       (Chlorine    monoxide    or]       TT^,^  i  -j 

C12O  —  j  .      \  —  HC1O  hypochlorous  acid, 

[ hypochlorous  anhydride  j 

r^in        /^ui    •      j-     -j        f  HC1O2  chlorous  acid. 
C1O2  -  Chlorine  dioxide  -    „.__  .       . , 

[  HClUs  chloric  acid 

C^OT  —  Perchloric  anhydride  —  HC1O4  perchloric  acid 

The  names  of  these  acids  should  be  learned  carefully  as  an 
illustration  of  the  principles  used  in  naming  acids.  For  chlorous 
and  chloric  acids  the  endings  correspond  to  those  which  are  used 
for  oxides  and  chlorides  (p.  29).  The  prefix  hypo-  means  under, 
and  the  prefix  per-  means  above  or  beyond.  The  endings  and 
prefixes  refer  to  the  relative  amounts  of  oxygen  for  different 
acids  of  the  same  element.  The  relations  for  other  elements  are 
not  always  so  simple.  Thus  the  acids  of  sulphur  are : 

Sulfurous  acid,  H2SOs 
Sulfuric  acid,  H2SO4 
Persulfuric  acid,  H2S2O8  [HSO4]2 

The  persulfuric  acid  is  in  a  higher  state  of  oxidation  than 
sulfuric  acid  because  it  contains  less  hydrogen,  not  because  it 
contains  more  oxygen  in  proportion  to  the  sulfur. 


124  A  TEXTBOOK  OF  CHEMISTRY 

The  salts  of  the  acids  are  named  by  changing  the  -OILS  of  the 
acid  to  -ite  for  the  salt,  and  the  -ic  of  the  acid  to  -ate  for  the  salt. 

Hypochlorous  acid,  HC1O,  gives  potassium  hypochlorite,  KC1O 
Chlorous  acid,  HC1O2,  gives  potassium  chlorite,         KC1O2 

Chloric  acid,  HC1O3,  gives  potassium  chlorate,        KC1O3 

Perchloric  acid,         HC1O4,  gives  potassium  perchlorate,  KC1O4 

Hypochlorous  Acid.  Hypochlorites.  When  chlorine  is  dis- 
solved in  water,  it  has  been  pointed  out  that  a  small  amount  of 

hypochlorous  acid  is  formed  by  the  reversible  reaction : 

i 

C12  +  HOH  ^±  HC1  +  HC1O 

The  equilibrium  in  this  reaction  is  very  far  toward  the  left, 
but  if  a  base  is  added  to  the  solution,  the  two  acids  will  be  neutral- 

HC1  +  KOH  =  KC1  +  H20 
HC1O  +  KOH  =  KC1O  +  H2O 

The  neutralization  of  the  acids  causes  a  displacement  of  the 
equilibrium  toward  the  right  side  of  the  first  equation  and  the 
reaction  goes  on  to  completion.  By  adding  the  three  equations 
together  and  eliminating  water,  the  result  can  be  expressed  in  the 
single  equation : 

2  KOH  +  C12  =  KC1  +  KOC1  +  H2O 

If  slaked  lime  (calcium  hydroxide,  Ca(OH)2)  is  used,  a  mixture 
of  calcium  chloride,  CaCl2,  and  calcium  hypochlorite,  Ca(OCl)2, 

,C\ 

or  a  calcium  chloride-hypochlorite,  Ca<^         ,  is  formed.     This 

X)C1 
is  called  bleaching  powder. 

/Cl 

Ca(OH)2  +  C12  =  Ca<        +  H2O 
•OCl 

Calcium 
Chloride-hypochlorite 


HYPOCHLORITES  125 

The  hypochlorites  give  up  their  oxygen  readily  to  other  sub- 
stances and  so  are  powerful  oxidizing  agents.  The  action  is 
much  more  vigorous  in  a  faintly  acid  than  in  an  akaline  solution, 
however,  because  hypochlorous  acid,  HC1O,  gives  up  its  oxygen 
much  more  easily  than  a  hypochlorite  does.  For  this  reason 
bleaching  powder  is  applied  to  the  bleaching  of  cotton  or  linen 
cloth  by  dipping  the  cloth  first  in  a  solution  of  the  bleaching 
powder  and  then  in  very  dilute  acid,  which  liberates  the  hypo- 
chlorous  acid. 

Not  only  may  a  hypochlorite  be  used  to  oxidize  other  sub- 
stances, but  if  a  neutral  or  faintly  acid  solution  of  a  hypochlorite 
is  boiled,  one  portion  is  oxidized  to  a  chlorate  while  another  por- 
tion is  reduced  to  a  chloride  : 

KC1O  +  2  KC1O  =  KC1O3  +  2  KC1 

Potassium 
Chlorate 

If  a  small  amount  of  a  cobalt  salt,  as  cobalt  nitrate,  Co(NOs)2, 
is  added  to  a  solution  of  a  hypochlorite,  the  oxygen  of  the  hypo- 
chlorite is  liberated  in  the  free  state.  The  cobalt  is  oxidized  to 
cobalt  dioxide,  CoO2,  which  then  acts  as  a  catalyzer,  as  man- 
ganese dioxide  acts  on  potassium  chlorate  : 

,C\ 

2  Ca<         +  CoO2  =  2  CaCl2  +  O2  +  CoO2 
XOC1 

Hypochlorous  acid  is  a  very  weak  acid.  While  in  the  reaction 
of  ionization  : 

HCI  ^t  H+  +  cr    . 

which  occurs  when  hydrochloric  acid  is  dissolved  in  water,  the 
equilibrium  is  far  to  the  right  in  moderately  dilute  solutions,  for 
hypochlorous  acid  the  corresponding  reaction  : 


has  the  equilibrium  very  far  to  the  left.  In  other  words  hydro- 
chloric acid  gives  a  large  proportion  of  hydrogen  ions  in  dilute 
solutions,  while  hypochlorous  acid  gives  only  a  very  small  pro- 


126  A  TEXTBOOK  OF   CHEMISTRY 

portion  of  such  ions.  This  fact  may  be  used  to  obtain  a  solu- 
tion of  hypochlorous  acid.  If  hydrochloric  acid  is  added  to  a 
solution  of  a  hypochlorite,  hypochlorous  acid  will  be  formed  in 
accordance  with  the  reaction  : 

K+  +  cio-  +  H+  +  cr  =  HCIO  +  K+  +  cr 

From  such  a  solution  hypochlorous  acid  and  water  pass  over 
together,  on  distillation,  and  this  is  the  easiest  method  of  getting 
a  solution  of  hypochlorous  acid.  An  excess  of  hydrochloric  acid 
must  be  avoided,  however,  as  this  would  cause  the  reaction : 

H+  +  Cl-  +  HCIO  ^±  H2O  +  C12 

to  occur,  in  which  the  equilibrium  is  far  to  the  right. 

Hypochlorous  acid  can  be  obtained  only  in  dilute  solutions. 
Concentrated  solutions  decompose  in  accordance  with  the  re- 
actions already  given : 

2  HCIO  =  2  HC1  +  O2 
HCIO  +  2  HCIO  =  HC103  +  2  HC1 
HC1  +  HCIO  =  H20  +  C12 

*  Hypochlorous  Anhydride  or  Chlorine  Monoxide.  Chlorine 
monoxide,  C12O,  is  formed  when  chlorine  is  passed  through  a 
tube  containing  cold,  dry  mercuric  oxide,  the  mercury  being  con- 
verted into  an  oxychloride : 

2  HgO  +  2  C12  =  HgO.HgCl,  +  C12O 

The  oxide  of  mercury  used  must  be  obtained  by  precipitation, 
and  washed  and  dried  at  300°  —  400°,  as  the  crystalline  oxide 
does  not  react  readily  enough.  Chlorine  monoxide  may  be  con- 
densed to  a  liquid  which  boils  at  about  5°.  Either  the  liquid 
or  the  gas  explodes  violently  on  slight  provocation,  giving  chlo- 
rine and  oxygen : 

2  C12O  =  2  C12  +  O2 

In  this  case  the  affinity  between  atoms  of  the  same  kind  is 
greater,  apparently,  than  that  between  chlorine  and  oxygen  in 
chlorine  monoxide.  Curiously  enough,  when  chlorine  combines 


CHLORATES  127 

with  a  larger  amount  of  oxygen,  the  compound  is  much  more 
stable. 

*  Chlorous  Acid  and  Chlorites.     When  sodium  peroxide  is 
added  to  a  solution  of  chlorine  peroxide,  C1O2,  sodium  chlorite 
is  formed : 

2  C1O2  +  Na2O2  =  2  NaClO2  +  O2 

The  chlorites  are  bleaching  agents,  similar  to  the  hypochlo- 
rites,  and  are  even  more  unstable.  Free  chlorous  acid  has  not 
been  prepared,  even  in  solution. 

Chloric  Acid  and  Chlorates.  When  a  faintly  acid  solution 
of  a  hypochlorite  is  warmed,  one  portion  oxidizes  another  to  a 
chlorate : 

KC1O  +  2  KC10  =  KC103  +  2  KC1 

Practically,  if  chlorine  is  passed  into  a  solution  of  potassium 
hydroxide  or  into  milk  of  lime,  Ca(OH)2,  till  there  is  a  slight 
excess  and  the  solution  becomes  warm  from  the  heat  evolved  by 
the  reaction,  a  solution  of  potassium  chloride  and  potassium 
chlorate,  or  of  calcium  chloride  and  calcium  chlorate  will  be  ob- 
tained. The  student  should  write  the  equations  and  notice 
what  portion  of  the  potassium  is  converted  into  potassium  chlo- 
rate. Why  is  it  more  economical  to  prepare  calcium  chlorate 
first  and  then  obtain  potassium  chlorate  by  adding  potassium 
chloride  to  the  solution  ?  What  must  be  the  relative  solubilities 
of  calcium  chlorate  and  potassium  chlorate  for  such  a  method 
to  be  successful  ? 

Chloric  acid  is  stable  only  in  solution  and  cannot  be  separated 
as  a  pure  compound. 

*  Chlorine    Dioxide.      When   concentrated   sulfuric   acid   is 
added  to  a  chlorate  the  chloric  acid  liberated  decomposes  at  once 
into  perchloric  acid,  chlorine  peroxide  and  water : 

KC1O3  +  H2S04  =  HC103  +  KHS04 

3  HC103  =  HC104  +  2  C102  +  H2O 

Chlorine  dioxide  is  a  heavy,  yellow  gas,  which  is  easily 
soluble  in  water.  It  seems  to  be  even  more  unstable  than  chlo- 


128  A  TEXTBOOK  OF   CHEMISTRY 

rine  monoxide  and  explodes  violently  if  warmed  or  brought  into 
contact  with  organic  matter.  These  properties  may  be  illus- 
trated by  mixing  some  sugar  and  potassium  chlorate  and  adding 
a  drop  of  concentrated  sulfuric  acid.  The  chlorine  dioxide 
will  react  with  the  sugar  and  ignite  the  mixture. 

The  solution  of  chlorine  peroxide  in  water  gives  with  a  base 
a  mixture  of  chlorite  and  chlorate : 

2  C1O2  +  2  KOH  =  KC1O2  +  KC103  +  H2O 

Chlorine  dioxide  may,  therefore,  be  considered  as  an  anhy- 
dride of  both  chloric  and  chlorous  acids. 

Perchlorates  and  Perchloric  Acid.  When  potassium  chlorate 
is  heated  to  its  melting  point,  it  partly  decomposes  into  potassium 
chloride  and  oxygen,  but  a  part  is  oxidized  to  potassium  per- 
chlorate,  while  another  part  is  reduced  to  potassium  chloride : 

3  KC1O3  +  KC1O3  =  3  KC1O4  +  KC1 

Potassium 
Perchlorate 

This  illustrates,  again,  the  fact  that  the  compounds  of  chlorine 
with  oxygen  become  more  and  more  stable  as  more  oxygen  is 
taken  up.  This  is  true  of  the  acids  as  well  as  of  the  salts.  Per- 
chloric acid  is  the  only  oxyacid  of  chlorine  which  can  be  ob- 
tained as  a  pure  compound,  free  from  water.  Perchloric  an- 
hydride, C^OT,  is  also  the  most  stable  of  the  oxides  of  chlorine. 

It  has  been  pointed  out  that  concentrated  sulfuric  acid  expels 
hydrochloric  acid  from  salt  chiefly  because  hydrochloric  acid 
is  a  gas  and  escapes  from  the  mixture ;  also  that  concentrated 
hydrochloric  acid  will  precipitate  salt  from  a  solution  of  acid 
sodium  sulfate,  NaHSO4,  leaving  sulfuric  acid  in  solution,  chiefly 
because  the  salt  is  nearly  insoluble  in  concentrated  hydrochloric 
acid.  Both  principles  may  be  used  to  prepare  perchloric  acid. 
If  concentrated  hydrochloric  acid  is  poured  over  some  sodium 
perchlorate  (30  cc.  for  20  grams  of  perchlorate),  the  reversible 
reaction : 

NaC104  +  HC1  =  NaCl  +  HC1O4 


PERCHLORIC  ACID  129 

will  proceed  till  95  per  cent  of  the  sodium  separates  as  sodium 
chloride.  This  may  be  removed  by  filtering  on  an  asbestos  filter 
and  washing  the  salt  with  a  small  amount  of  concentrated  hy- 
drochloric acid.  The  filtrate1  contains  a  little  salt  with  hy- 
drochloric and  perchloric  acids.  The  highest  boiling  point  of 
an  aqueous  solution  of  hydrochloric  acid  is  110°,  while  the  boil- 
ing point  of  the  hydrated  perchloric  acid  is  203°.  On  heating 
the  mixture,  therefore,  the  hydrochloric  acid  escapes  and  finally 
the  reversible  reaction : 

NaCl  +  HC104  =  NaClO4  +  HC1 

is  carried  to  completion,  leaving  only  perchloric  acid  containing 
a  small  amount  of  sodium  perchlorate.  A  pure  hydrated  per- 
chloric acid,  containing  about  28  per  cent  of  water,  may  be  ob- 
tained from  the  mixture  by  distilling  under  diminished  pressure. 
This  hydrated  acid  boils  with  some  decomposition  at  203°,  under 
atmospheric  pressure.  Anhydrous  perchloric  acid  may  be  ob- 
tained by  distilling  a  mixture  of  potassium  perchlorate  and  con- 
centrated sulfuric  acid  under  diminished  pressure.  The  anhy- 
drous acid  is  far  less  stable  than  the  hydrated  acid.  This  is 
possibly  because  the  hydrated  acid  has  the  structure : 

/0-H 


:%o 

^O 

*  Perchloric  Anhydride,  C12O7,  may  be  prepared  by  adding 
perchloric  acid  to  phosphoric  anhydride  cooled  to  10°  and  dis- 
tilling after  some  time : 

2  HC1O4  +  P2O5  =  C1207  +  2  HPO3 

Metaphosphoric 
Acid 

Perchloric  anhydride  is  a  colorless,  oily  liquid  which  boils  at 

82°. 

1  The  portion  of  a  solution  which  has  passed  through  a  filter. 


130  A  TEXTBOOK  OF  CHEMISTRY 

Structure  of  the  Oxyacids  of  Chlorine.  If  chlorine  is  univa- 
lent  in  the  oxyacids  of  chlorine,  the  structure  of  these  acids  would 
be  represented  by  the  formulas  : 

H  —  O  —  Cl  Hypochlorous  acid 

H—  O—  O—  Cl  Chlorous  acid 

H—  O—  O—  O—  Cl         Chloric  acid 
H—  O—  O—  O—  O—  Cl  Perchloric  acid 

There  is  very  little  evidence  that  oxygen  atoms  can  unite  in 
this  manner  to  form  chains  that  are  stable,  and  the  instability 
of  hydrogen  peroxide,  H  —  O  —  O  —  H,  makes  these  formulas 
seem  very  improbable.  There  is  also  a  good  deal  of  evidence 
to  show  that  chlorine  and  other  related  elements  may  have  a 
valence  as  high  as  seven  in  some  of  their  compounds.  The  fol- 
lowing structures  are,  therefore,  considered  much  more  probable  : 

^O 
H—  O—  Cl  H—  O—  C1=O        H—  O—  Cl^O 

H—  O—  ClO 


These  compounds  illustrate  very  clearly  the  effect  of  oxygen 
in  giving  an  acid  character  to  hydrogen  compounds.  Hypo- 
chlorous  acid  is  a  very  weak  acid  and  separates  almost  as  easily 
into  chlorine  and  hydroxyl,  H  —  O  —  ,  as  it  does  into  hydrogen 
and  —  O  —  Cl.  Perchloric  acid  on  the  contrary,  is  a  strong, 
stable  acid,  especially  in  the  hydrated  form.  For  a  possible 
explanation  see  p.  206. 

*  The  Atomic  Weight  of  Chlorine.  As  oxygen  is  the  basis  for 
atomic  weights,  it  would  be  most  natural  to  determine  the  atomic 
weight  of  chlorine  by  determining  the  composition  of  one  of  the 
oxides  of  chlorine.  But  these  oxides  are  so  unstable  that  they 
cannot  be  prepared  in  a  condition  of  sufficient  purity  for  such 
a  purpose.  The  composition  of  hydrochloric  acid  has  been  de- 
termined accurately,  however,  in  two  ways  ;  first,  by  combining 
a  weighed  amount  of  hydrogen  with  chlorine,  which  was  weighed 
in  the  liquid  form;  second,  by  passing  a  weighed  amount  of 


ATOMIC  WEIGHT   OF  CHLORINE  131 

hydrogen  over  potassium  chloroplatinate,  K2PtCl6,  which  was 
reduced  to  potassium  chloride  and  metallic  platinum.  The 
loss  in  weight  gave  the  weight  of  the  chlorine,  while  the  hydro- 
chloric acid  formed  was  also  collected  and  weighed.  The  aver- 
age for  the  ratio  between  hydrogen  and  chlorine  by  the  two 
methods  is,  H :  Cl  =  1 :  35.189.  As  the  atomic  weight  of  hy- 
drogen is  1.0078  (p.  72),  the  atomic  weight  of  chlorine  is 

35.189  X  1.0078  =  35.463 

The  atomic  weight  has  also  been  very  accurately  determined 
by  an  entirely  different  method.  By  dissolving  lithium  chloride, 
LiCl,  in  a  solution  of  perchloric  acid  and  evaporating  the  water, 
it  was  converted  into  lithium  perchlorate,  LiClO4.  The  ratio  of 
the  increase  in  weight  to  the  weight  of  the  lithium  chloride  was  : 

4  0  :  LiCl  =  1.50968  :  1  =  64  :  42.393 

This  gives  the  molecular  weight  of  LiCl  as  42.393. 

Next,  the  amount  of  silver  required  to  combine  with  the  chlo- 
rine of  the  lithium  chloride  was  determined  by  dissolving  a 
weighed  amount  of  silver  in  nitric  acid  and  adding  the  solution 
of  silver  nitrate,  AgNOs,  to  the  solution  of  lithium  chloride. 
This  gave  the  ratio  : 

LiCl :  Ag  =  0.39299  :  1  =  42.393  :  107.871 

This  gives  the  atomic  weight  of  silver,  Ag  =  107.871.  The 
amount  of  silver  chloride  which  could  be  obtained  from  a  given 
weight  of  lithium  chloride  was  also  determined.  This  gave : 

LiCl :  AgCl  =  0.295786 : 1  =  42.293  :  143.325 

Subtracting  the  atomic  weight  of  silver  from  the  molecular 
weight  of  silver  chloride  we  have  : 

143.325  -  107.871  =  35.454 

which  is  the  atomic  weight  of  chlorine.  It  will  be  seen  that  this 
value  agrees  cjosely  with  that  given  by  the  other  methods,  and 
it  does  not  seem  likely  that  the  value  35.46,  which  is  given  in 
the  atomic  weight  tables,  can  be  far  wrong. 


CHAPTER  IX 

CLASSIFICATION   OF   THE   ELEMENTS.     THE   PERIODIC 

SYSTEM 

THE  three  elements  which  have  been  studied,  oxygen,  hydrogen 
and  chlorine,  differ  very  greatly  from  each  other,  but  as  we  pass 
on  to  other  elements,  it  will  be  found  that  several  of  these  have 
very  marked  resemblances  to  chlorine,  while  others  have  proper- 
ties which  recall  those  of  oxygen,  though  the  resemblances  are 
not  so  close.  The  elements  fall  into  a  number  of  more  or  less 
well-defined  families  or  groups,  and  a  knowledge  of  these  groups  is 
of  great  assistance  in  acquiring  a  knowledge  of  the  elements  and 
their  compounds.  The  most  satisfactory  classification  is  the  one 
known  as  the  Periodic  System,  which  is  based  on  the  atomic 
weights  and  is  found  in  the  accompanying  tables.  The  elements 
are  arranged  in  the  order  of  their  atomic  weights,  with  a  few 
exceptions,  which  will  be  referred  to  below.  Hydrogen  does 
not  seem  to  fall  into  the  classification  and  is  omitted  in  the  first 
table.  The  valence  of  hydrogen  would  put  it  in  Group  I,  while 
its  amphoteric  (p.  206)  character  in  HOH  would  relate  it  to 
aluminium  or  silicon.  Beginning  with  helium,  the  first  seven 
elements  after  helium,  He,  4;  Li,  6.94;  Be,  9;  B,  11 ;  C,  12 ; 
N,  14;  O,  16;  F,  19,  pass  from  lithium,  which  is  strongly 
metallic,  to  fluorine,  which  is  very  strongly  nonmetallic.  The 
first  five  form  oxygen  compounds  as  follows :  L^O,  BeO,  B2O3, 
CO2,  N2Os.  The  last  four  form  compounds  with  hydrogen, 
CH4,  NH3,  OH2,  FH. 

In  the  second  row  we  find  that  the  ninth  element,  neon, 
resembles  helium,  sodium  resembles  lithium,  magnesium  re- 
sembles beryllium,  and  so  on  to  chlorine,  which  resembles 
fluorine.  Compounds  with  oxygen  and  hydrogen  are  : 

132 


CLASSIFICATION  OF  THE   ELEMENTS  133 

Na2O,  MgO,  A1203,  SiO2,  P2(V,  SO3,  C12O7 
—       —       —     SiH4,  PH3,  SH2,  C1H 

In  the  third  row  the  oxides  are  :  K2O,  CaO,  Sc2O3,  TiO2,  V2O5, 
CrO3,  Mn2O7,  but  the  last  four  elements  do  not,  as  in  the  first  two 
rows,  form  compounds  with  hydrogen.  At  the  end  of  the  row  are 
three  elements,  iron,  cobalt  and  nickel,  which  resemble  manganese 
in  some  of  their  properties.  Similar  groups  of  three  elements  are 
found  after  the  fifth  and  seventh  rows  of  elements.  Beginning 
with  the  third  row,  the  elements  of  alternate  rows  resemble  each 
other  much  more  closely  than  those  of  the  successive  rows,  and 
each  pair  of  rows  taken  together  is  spoken  of  as  a  long  period 
.to  distinguish  these  from  the  short  periods  of  the  first  two  rows. 

It  will  be  noticed  that  the  highest  valence  of  the  elements 
toward  oxygen  increases  from  "left  to  right,  from  one  to  seven  : 


Bf  M     N^O         x,0     C\fO 


fl 

"\  "\  J™  --V        ^  x^«  -.         =O 

)0,     Be=0,         V),     Cf    ,        >0,     S=0,         >0. 
LiX  B<  ^O     N^O       %     Cl%0 

^O  ?0 


The  valence  toward  hydrogen,  however,  decreases  from  the 
center  to  the  right  from  four  to  one  : 

H 

I  /H        /H 

H— C— H,     Nf-H,    O<     ,     F-H. 

T  \H  H 

H 

If  a  line  is  drawn  in  the  table  between  beryllium  and  boron 
downward  to  the  right  and  between  tellurium  and  tungsten  (W), 
all  of  the  elements  below  and  to  the  left  of  the  line  will  be  found 
to  be  metallic,  except  those  of  the  argon  family,  while  in  the  third, 
fifth  and  seventh  rows  the  elements  to  the  right  of  the  line  are 
also  metallic.  The  remaining  elements  are  nonmetallic.  These 

1  The  true  formula  is  P4Oi0. 


134 


A  TEXTBOOK  OF  CHEMISTRY 


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S>2*^ 


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2:      JH^ 

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***  CO 


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CLASSIFICATION  OF  THE  ELEMENTS  135 


S3 


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o  J2 


is 


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136  A  TEXTBOOK  OF  CHEMISTRY 

facts  and  also  the  short  and  long  periods  are  better  shown  in 
the  form  of  the  table  on  p.  135.  In  this  second  table  the  melting 
points  of  the  elements  are  also  given  and  it  is  very  clear  that 
these  increase  to  a  maximum  in  the  fourth  to  the  sixth  groups 
of  the  longer  periods  (horizontal  rows)  and  fall  off  on  either  side. 
This  periodicity  of  the  melting  points  has  proved  useful  in  indi- 
cating those  elements  which  are  suitable  for  the  filaments  of  in- 
candescent lights. 

There  is  also  a  periodic  relation  between  the  atomic  weights 
of  the  elements  and  their  atomic  volumes.  The  atomic  volume 
may  be  defined  as  the  volume  occupied  by  one  gram-atom  of  the 
element.  Thus,  the  specific  gravity  of  potassium  is  0.862  and  the 

39  1 
atomic  volume  is       '      =  45.4.      The  specific  gravity  of  silver 


is  10.492  and  the  atomic  volume  is  ±^^  =  W.2S.     The  peri- 

odic relation  between  atomic  weights  and  atomic  volumes  is 
clearly  shown  in  Fig.  46. 

When  the  Periodic  Table  was  first  proposed  by  Mendeleef, 
scandium,  germanium  and  several  other  elements,  which  have 
been  discovered  since  then,  were  unknown,  and  he  predicted  the 
discovery  of  these  elements  and  pointed  out  some  of  their  prop- 
erties. When  these  elements  were  discovered  a  few  years  later, 
the  fulfillment  of  this  prophecy  helped  very  much  toward  the 
acceptance  of  the  table  among  chemists.  Still  later  the  table 
caused  Professor  Ramsay,  after  the  discovery  of  argon  and 
helium,  to  search  diligently  for  the  other  elements  of  the  Zero 
Group  which  were  predicted  by  the  table.  His  search  proved 
successful  and  resulted  within  a  few  years  in  the  discovery  of 
neon,  krypton,  xenon  and  niton.  The  group  is  called  the  zero 
group  because  the  elements  of  the  group  do  not  combine  with 
other  elements  and  their  valence  is  considered  to  be  zero. 

The  atomic  weight  which  has  been  accepted  for  tellurium  is 
greater  than  it  should  be  in  accordance  with  the  properties  of 
the  element,  which  place  it  in  the  sulfur  family.  This  has  led 
a  number  of  investigators  to  examine  the  compounds  of  the  ele- 


CLASSIFICATION  OF  THE  ELEMENTS  137 


OIHOIV 


138  A  TEXTBOOK  OF  CHEMISTRY 

ment  very  carefully  and  to  redetermine  the  atomic  weight  by 
different  methods. 

In  three  other  cases  (A  and  K ;  Co  and  Ni ;  Pr  and  Nd)  the 
relative  positions  of  the  elements  do  not  correspond  to  the  prop- 
erties, and  all  of  these  elements  have  been  very  carefully  studied 
for  this  reason.  No  one  has  been  able  to  show,  however,  that 
the  accepted  atomic  weights  for  these  eight  elements  are 
wrong,  and  we  are  forced  to  the  conclusion  that  the  same  factors 
which  cause  the  differences  between  successive  atomic  weights  to 
be  irregular  have,  in  these  cases,  displaced  the  elements  from 
what  seem  to  be  their  normal  places. 

The  rare  elements  given  in  the  footnote  on  p.  134  are  not  easily 
placed  in  the  ordinary  forms  of  the  Periodic  Table ;  Werner  and 
others  have  proposed  arrangements  which  include  these  elements 
on  the  general  principle  that  as  the  short  periods,  He — F  and 
Ne — Cl  are  followed  by  longer  periods,  A — Br  and  Kr — I, 
these  periods  are,  in  turn,  followed  by  periods  containing  each 
a  still  larger  number  of  elements. 

It  is  clear  from  what  has  been  said  that  the  Periodic  System 
is  not  only  useful  as  a  convenient  means  of  classifying  the  ele- 
ments and  for  didactic  purposes,  but  that  it  has  also  proved 
a  powerful  stimulus  to  chemical  research. 

The  relations  between  the  atomic  weights  of  the  elements  and 
their  properties  which  are  brought  out  in  the  Periodic  System 
constantly  suggest  that  the  elements  must  have  some  common 
origin  and  that  the  atoms  are  complex  aggregates  built  up  in 
some  way  from  simpler  parts.  Such  an  idea  has  received  very 
strong  support  from  the  phenomena  connected  with  radium  and 
other  radioactive  elements  (p.  471). 


CHAPTER   X 
THE  HALOGEN  FAMILY 

General  Properties  of  the  Halogens.  The  four  elements  of 
the  halogen  family  are :  x 

Fluorine,  F,  19 

Chlorine,  Cl,  35.5 

Bromine,  Br,  80 

Iodine,  I,  127 

The  elements  of  the  halogen  group  are  the  most  strongly  non- 
metallic  of  all  the  elements.  They  are  also  called  negative  be- 
cause in  the  electrolysis  of  their  compounds  they  are  attracted 
toward  the  anode  or  positive  electrode.  In  contrast  with  these 
and  other  nonmetallic  elements,  metals  are  called  positive,  the 
most  strongly  positive  or  metallic  elements  being  those  of  the 
alkali  group,  to  which  sodium  and  potassium  belong.  The  name 
halogen  means  "  salt-former,"  and  is  given  to  these  elements  be- 
cause they  combine  directly  with  metals  to  form  salts,  sodium 
chloride  or  common  salt,  NaCl,  being  the  most  important  ex- 
ample. The  most  common  salts  containing  other  nonmetallic 
elements  are  those  which  also  contain  oxygen,  as  sodium  sulfate, 
Na2SO4,  or  potassium  nitrate,  KNO3. 

Compounds  of  the  Halogens  with  Hydrogen  and  Oxygen. 
The  compounds  of  the  halogens  with  hydroge'n  and  with  hydro- 
gen and  oxygen  are  acids  and  have  the  following  formulas : 

H2F2orHF          HC1  HBr  HI 

HC10  HBrO  HIO 

HC102 

HC103  HBrO3  HIO3 

HC1O4  HI04 

1  In  this  and  other  similar  tables  approximate  values  are  given 
for  the  atomic  weights  in  order  that  the  student  may  learn  the  rela- 

139 


140  A  TEXTBOOK  OF  CHEMISTRY 

The  elements  of  the  group  are  univalent  in  combining  with 
hydrogen  or  with  positive  elements  and  frequently,  also,  in  com- 
bining with  nonmetallic  elements.  In  combining  with  oxygen 
or  with  oxygen  and  hydrogen  the  valence  seems  to  vary  from 
one  in  hypochlorous  acid,  H  —  O  —  Cl,  to  seven  in  perchloric 


acid,  H  —  O  —  Cl^O,  the  odd  numbers  of  valences  being  most 


common. 

Fluorine  is  the  most  strongly  nonmetallic  or  negative  element 
of  the  group,  or,  indeed,  of  all  of  the  elements.  It  will  displace 
any  other  element  of  the  group  from  combination  with  hydrogen 
or  a  metal.  In  a  similar  manner  chlorine  will  displace  bromine 
or  iodine,  and  bromine  will  displace  iodine.  This  is  prob- 
ably due  to  the  same  properties  which  cause  hydrofluoric 
acid  to  be  the  most  stable  and  hydriodic  acid  to  be  the  least 
stable  of  'the  compounds  of  these  elements  with  hydrogen. 

The  resemblances  between  chlorine,  bromine  and  iodine  are 
much  closer  than  the  resemblances  between  these  elements  and 
fluorine.  For  this  reason  bromine  and  iodine  are  considered 
first.  In  studying  these  elements  the  properties  of  chlorine 
and  its  compounds  should  be  constantly  recalled  and  the  re- 
semblances emphasized. 

Bromine,  Br,  79.92.  Occurrence,  Preparation.  In  most  cases 
where  large  quantities  of  chlorides  are  found  in  nature  smaller 
amounts  of  bromides  are  found  associated  with  them.  In  this 
way  bromides  are  found  especially  in  sea  water  and  in  the  brines 
from  which  salt  is  obtained  by  evaporation  and  crystallization. 
Some  of  the  American  brines  in  Michigan  are  rich  in  bromine, 
and  the  bromine  is  obtained  from  these  by  subjecting  them  to 
electrolysis  till  all  of  the  bromine  is  liberated,  with  a  small 
amount  of  chlorine.  As  bromine  boils  at  59°  and  has  a  molec- 
ular weight  of  160,  on  boiling  the  liquid  the  bromine  will  pass 
off  with  a  comparatively  small  amount  of  water.  (What  con- 

tions  among  the  atomic  weights  more  easily.     Accurate  values  are 
given  on  p.  10. 


BROMINE 


141 


nection  has  the  last  fact  with  the  molecular  weight  of  bromine  ?) 
The  chlorine  may  be  removed  by  mixing  the  impure  bromine 
with  a  solution  of  potassium  bromide  and  distilling  : 

2  KBr  +  C12  =  2  KC1  +  Br2 

Bromine  may  also  be  prepared  by  warming  a  mixture  of  potas- 
sium bromide,  manganese  dioxide  and  sulfuric  acid.  The  prod- 
ucts are  potassium  sulfate,  manganese  sulfate,  bromine  and 
water.  What  is  the  equation  for  the  reaction? 

Properties.  Bromine  is  a  heavy,  very  dark  colored  liquid, 
which  gives  off  reddish  brown  vapors  at  ordinary  temperatures. 
It  has  a  strong,  disagreeable 
odor,  the  name  having  been 
given  to  it  for  this  reason, 
from  /Spw/Aos,  a  stench.  It 
is  also  an  irritant  poison. 
As  with  chlorine,  the  best 
antidote  is  to  breathe  the 
vapor  of  strong  alcohol. 
If  the  liquid  touches  the 
skin,  it  produces  a  severe 
wound,  which  it  is  very 
difficult  to  heal. 

Although  bromine  vapor  is  much  heavier  than  air  (how  many 
times  heavier  ?),  if  a  little  of  the  liquid  is  placed  in  the  bottom  of 
a  tall  cylinder,  the  vapor  will  diffuse  rapidly  upward  through  the 
air  in  the  cylinder.  How  can  this  be  explained  by  the  kinetic 
theory  ? 

Bromine  combines  directly  with  both  metals  and  nonmetals, 
forming  compounds  which  are,  in  almost  all  cases,  very  similar 
to  the  corresponding  chlorides  both  in  formulas  and  in  properties. 

At  228°  the  volume  of  bromine  vapor  which  would  fill  22.4 
liters  at  0°  and  760  mm.1  weighs  about  160  grams,  but  at  1570° 
the  gram  molecular  volume  weighs  only  a  little  over  100  grams. 


Fig.  47 


1  Supposing  that  the  vapor  could  be  cooled  to  this  temperature 
at  a  pressure  of  760  mm.  without  its  condensing  to  a  liquid. 


142  A  TEXTBOOK  OF  CHEMISTRY 

This  indicates  that  at  high  temperatures  bromine  is  largely 
dissociated  into  molecules  which  contain  only  a  single  atom. 
At  lower  temperatures  the  formula  of  bromine  is  evidently  Br2. 

Bromine  melts  at  —  7°  and  boils  at  59°.  It  has  a  specific 
gravity  of  3.1883  at  0°  or  of  2.9483  at  59°.  It  forms  a  hydrate 
which  probably  has  the  composition  Br2  +  8  H2O,  though  the 
analyses  of  the  compound  do  not  agree  very  well  with  the  for- 
mula. Potassium  and  sodium  bromides  are  used  in  medicine  as 
sedatives,  the  latter  by  preference,  because  the  bromide  ion  seems 
to  be  the  constituent  which  produces  the  desired  effect,  while  the 
potassium  ion  is  much  more  irritant  than  the  sodium  ion  when 
taken  in  moderate  quantities.  Silver  bromide  is  used  in  pho- 
tography, especially  in  the  preparation  of  "  dry  plates."  Many 
compounds  of  bromine  are  used  in  the  manufacture  of  coal- 
tar  dyes.  Bromine  has  also  been  used  to  a  limited  extent  as  a 
disinfectant. 

Hydrobromic  Acid.  From  the  method  used  in  preparing  hy- 
drochloric acid  we  should  expect  to  get  hydrobromic  acid  by  the 
action  of  sulfuric  acid  on  sodium  bromide  or  potassium  bromide : 

KBr  +  H2SO4  :£  KHSO4  +  HBr 

This  reaction  takes  place  when  the  substances  are  mixed,  but 
the  hydrobromic  acid  gas  which  escapes  will  be  colored  brown, 
indicating  the  presence  of  free  bromine.  Sulfur  dioxide,  SO2, 
is  also  found  in  the  vapor : 

H2SO4  +  2  HBr  ^  H2SO3    +  Br2  +  H2O 

Sulfurous 
Acid 

The  sulfurous  acid  is  unstable  and  decomposes  into  sulfur 
dioxide  and  water : 

H2SO3  ^1  H2O  +  SO2 

This  is  evidently  because,  owing  to  the  comparatively  weak 
affinity  between  bromine  and  hydrogen,  hydrobromic  acid  acts 
as  a  reducing  agent  toward  sulfuric  acid. 


BROMINE 


143 


To  obtain  hydrobromic  acid  free  from  bromine  a  mixture  of 
hydrogen  and  bromine  vapor  may  be  passed  through  a  tube  con- 
taining a  red-hot  spiral  of  platinum  wire. 

Another  method  is  to  drop  bromine  into  a  mixture  of  red  phos- 
phorus and  water  and  pass  the  hydrobromic  acid  gas  through  a 
tube  containing  red  phosphorus  and  glass  wool  moistened  with 
a  strong  solution  of  hydrobromic  acid.  The  method  depends  on 
the  hydrolysis  of  phosphorus  tribromide  by  water  (p.  115). 

2P     +3Br2       =     2PBr3 

PBr3  +  3  HOH  =     P(OH)3    +  3  HBr 

Phosphorous 
Acid 

Hydrobromic  acid  is  a  colorless  gas,  which  fumes  strongly  in 
the  air  owing  to  its  condensation  with  the  moisture  of  the  air 
to  form  a  concen- 
trated solution,  which 
has  a  much  lower 
vapor  pressure  than 
that  of  water.  Water 
dissolves  the  acid  even 
more  readily  than  it 
dissolves  hydrochloric 
acid.  The  constant 
boiling  mixture  of  hy- 
drobromic acid  and 
water  boils  at  125° 
and  contains  47.7  per 
cent  of  hydrobromic 
acid;  while  the  cor- 
responding mixture  of  hydrochloric  acid  and  water  boils  at  110° 
and  contains  only  20.24  per  cent  of  hydrochloric  acid.  The 
density  of  the  hydrobromic  acid  solution  is  also  considerably 
greater  for  a  given  per  cent  of  acid. 

Sodium  Hypobromite,  NaBrO,  is  obtained  by  dissolving  bro- 
mine in  a  cold  solution  of  sodium  hydroxide,  or,  better,  by  draw- 
ing the  vapor  of  bromine  through  the  solution  with  a  current 


Fig.  48 


144  A  TEXTBOOK  OF  CHEMISTRY 

of  air.     (See  hypochlorites.)       If  the  solution  is  warmed,  the 
hypobromite  is  changed  to  the  bromate,  NaBrO3 : 

2  NaBrO  +  NaBrO  =  2  NaBr  +  NaBrO3 

Iodine,  I,  126.92.  Occurrence,  Preparation.  While  a  minute 
quantity  of  iodine  is  found  in  sea  water  and  in  almost  all  brines, 
the  amount  is  too  small  for  the  practical  preparation  of  the  ele- 
ment. Many  seaweeds,  however,  absorb  a  small  amount  of 
iodine  from  the  sea  water.  The  ash  from  these  weeds  is  called 
kelp  and  contains  a  small  amount  of  iodides.  From  these  the 
iodine  may  be  liberated  by  chlorine  or  by  sulfuric  acid  and 
manganese  dioxide.  (If  sodium  iodide  is  the  compound  pres- 
ent, what  will  be  the  equation  for  the  required  action  ?) 

Iodine  is  also  found  as  sodium  iodate,  NaIO3,  in  the  crude 
sodium  nitrate  from  Chile  and  Peru.  The  crude  nitrate 
("  caliche  ")  contains  about  0.2  per  cent  of  this  compound  and 
most  of  the  iodine  of  commerce  comes  from  this  source. 

*  Iodine  is  found  in  the  thyroid  gland,  and  its  presence  seems 
to  be  physiologically  important.  The  diseases  of  goiter  and 
cretinism  are,  apparently,  connected  with  a  deficiency  of  iodine. 

Properties  of  Iodine.  Iodine  is  obtained  in  the  form  of 
black,  crystalline  scales  which  melt  at  114.2°.  The  liquid  boils 
at  184.3°,  but  gives  off  a  beautiful  violet  vapor  at  much  lower 
temperatures.  The  weight  of  the  vapor  indicates  that  the  for- 
mula is  I2  at  temperatures  not  far  above  the  boiling  point,  but 
even  at  700°  the  molecules  dissociate  appreciably  into  single 
atoms,  just  as  the  bromine  and  chlorine  molecules  dissociate 
at  much  higher  temperatures.  The  stability  of  iodides  is  much 
less  than  that  of  bromides  or  chlorides,  and  the  stability  of  the 
iodine  molecule  is  also  much  less  than  that  of  the  bromine  mole- 
cule. In  general  the  affinity  of  the  nonmetallic  elements  toward 
metallic  elements  decreases  with  increasing  atomic  weight. 

Iodine  dissolves  very  slightly  in  pure  water.  It  dissolves 
more  easily  in  alcohol,  giving  a  brown  solution,  called  tincture l  of 

1  The  name  tincture  is  given  in  pharmacy  to  a  solution  in  alcohol 
of  some  substance  or  of  the  active  constituents  of  some  plant. 


IODINE  145 

iodine,  which  is  used  in  medicine.  Iodine  also  dissolves  in  a  so- 
lution of  potassium  iodide.  There  is  evidence  that  in  solution  it 
forms  an  unstable  compound,  KIs,  but  the  dilution  of  a  solution 
having  this  composition  causes  the  precipitation  of  a  part  of  the 
iodine.  In  chloroform,  carbon  bisulphide  and  other  solvents 
with  which  it  does  not  combine,  iodine  forms  violet  solutions. 

Iodine  gives  with  starch  emulsion,  in  the  presence  of  hy- 
driodic  acid  or  an  iodide,  a  deep  blue  color,  which  is  very  charac- 
teristic and  which  is  used  as  a  test  for  free  iodine  or  for  starch. 
From  iodides  the  iodine  must  be  liberated  by  some  oxidizing 
agent,  best  by  nitrous  acid  for  the  detection  of  minute  quantities 
of  the  element.  Chlorine  may  be  used,  but  an  excess  oxidizes 
the  iodine  to  iodic  acid  and  destroys  the  color.  The  color  of 
the  starch  iodide  is  also  destroyed  by  heat,  but  returns,  in  part, 
on  cooling  the  solution. 

Hydriodic  Acid.  It  has  already  been  pointed  out  that  the 
affinity  between  hydrogen  and  iodine  is  much  less  than  that 
between  hydrogen  and  chlorine  or  bromine.  If  slightly  diluted 
sulfuric  acid  is  poured  on  potassium  iodide,  some  hydriodic  acid 
is  liberated,  but  the  larger  part  of  the  acid  acts  upon  more  of  the 
sulfuric  acid,  reducing  it  to  sulfur  dioxide  or  even  to  hydrogen 
sulfide,  H2S.  The  student  should  write  the  equations  for  the 
three  reactions  involved  in  the  last  statement  and  compare 
with  the  somewhat  similar  conduct  of  hydrobromic  acid. 

Hydriodic  acid  is  best  prepared  by  melting  together  in  a 
distilling  bulb  1  part  of  red  phosphorus  with  20  parts  of  iodine, 
forming  a  mixture  of  phosphorus  triodide,  Pis,  and  iodine.1 
When  the  mixture  is  cold,  a  stopper  bearing  a*  separatory  funnel 
containing  4  parts  of  water  is  fitted  to  the  neck  of  the  bulb,  and 
the  side  tube  is  connected  with  a  small  U-tube  containing  a  very 
little  water  to  wash  the  gas  and  retain  nearly  all  of  the  iodine 

1  Method  of  Lothar  Meyer  slightly  modified,  Ber.  20,  3381. 
The  usual  direction,  which  gives  enough  phosphorus  to  form  PIa» 
gives  rise  to  the  formation  of  phosphonium  iodide,  PH4I,  and  this 
may  stop  the  exit  tubes  or  contaminate  the  product.  For  the  same 
reason  iodine  cannot  well  be  removed  from  the  gas  by  moistened  red 
phosphorus,  as  directed  by  some  authors. 


146 


A  TEXTBOOK  OF  CHEMISTRY 


which  passes  over.  If  a  solution  of  hydriodic  acid  in  water  is 
desired,  the  delivery  tube  should  not  dip  beneath  the  surface  of 
the  water  which  is  to  absorb  the  gas. 

When  all  is  ready,  the  water  is  dropped  slowly  on  the  mixture 
of  phosphorus  triodide  and  iodine.  After  all  has  been  added 
the  last  of  the  hydriodic  acid  is  driven  over  by  warming.  The 
equation  is : 

PI3  +  I2  +  4  H2O  =  H3P04  +  5  HI 

Phosphoric 
Acid 

The  solution  of  hydriodic  acid  in  water  has  properties  similar  to 
those  of  the  corresponding  solutions  of  hydrochloric  and  hydrobro- 

mic  acids.  The  solu- 
tion of  constant  boiling 
point  under  atmos- 
pheric pressure  boils 
at  127°,  has  a  specific 
gravity  of  1.70  and 
contains  57  per  cent 
of  the  acid.  The  hy- 
driodic acid  is  slowly 
oxidized  with  libera- 
tion of  iodine  on  ex- 
posure of  the  solution 
to  the  air,  and  for  this 
reason  the  aqueous 
Fig.  49  acid  is  almost  always 

colored  red  or  brown. 

Direct  Combination  of  Hydrogen  and  Iodine.  Reversible  re- 
actions. Equilibrium.  Mixtures  of  equal  volumes  of  chlorine 
and  hydrogen  or  of  bromine  vapor  and  hydrogen  combine  com- 
pletely when  heated  to  the  temperature  of  combination,  and 
neither  hydrochloric  acid  nor  hydrobromic  acid  dissociates  ap- 
preciably unless  heated  to  a  quite  high  temperature.1  If  a  mix- 

1  Haber  (Thermodynaimk  technischer  Gasreaktionen,  1905,  S. 
95)  calculates  the  dissociation  of  hydrobromic  acid  as  only  0.15  per 
cent  at  727°. 


EQUILIBRIUM 


147 


ture  of  equal  volumes  of  hydrogen  and  of  iodine  vapor  is  heated, 
however,  there  is  no  temperature  at  which  the  combination  to 
form  hydriodic  acid  will  be  complete,  even  if  the  mixture  is 
heated  for  an  indefinitely  long  time.  On  the  other  hand,  if 
hydriodic  acid  is  heated,  it  decomposes  slowly,  even  at  quite 
low  temperatures,  but  never  completely,  no  matter  how  long  it 
is  heated,  unless  the  temperature  is  very  high  indeed.  If  we 
start  with  two  sealed  glass  tubes,  one  containing  one  part  by 
weight  of  hydrogen  with  127  parts  of  iodine  and  the  other 
containing  hydriodic  acid,  and  heat  both  tubes  at  the  same  tem- 
perature till  the  composition  no  longer  changes,  it  will  be  found 
that  each  tube  contains,  on  cooling,  a  mixture  of  hydrogen, 
iodine  and  hydriodic  acid  but  that  the  composition  of  the  mix- 
ture in  the  two  tubes  is  identical.  This  result  is  most  easily 
explained  by  supposing  that  we  have  here  a  reversible  reaction  : 

H2  +  I2  ^±  HI  +  HI 

and  that  when  equilibrium  is  reached  the  reaction  does  not  stop 
but  continues  in  such  a  way  that  just  as  many  molecules  of  hy- 
driodic acid  are  formed  in  a  minute  as  are  decomposed  in  the 
same  time. 

The  composition  of  the  mixture  which  is  in  equilibrium 
varies  with  the  temperature,  as  will  be  seen  from  the  following 
table : 


COMPOSITION  OF  THE  EQUILIBRIUM  MIXTURE  OF  HYDROGEN, 
IODINE  AND  HYDRIODIC  ACID 


TEMPERATURE 

PROPORTION  OP 
HYDRIODIC  ACID 

PROPORTION  OF  HYDROGEN 
AND  IODINE 

283° 

0.8213 

0.1787 

328° 

0.8115 

0.1885 

374° 

0.7990 

0.2010 

393° 

0.7942 

0.2058 

427° 

0.7843 

0.2157 

508° 

0.7592 

0.2408 

148 


A  TEXTBOOK  OF  CHEMISTRY 


It  is  evident  from  this  table  that  when  equilibrium  is  reached, 
the  combination  of  hydrogen  and  iodine  is  more  nearly  complete 
at  low  temperatures  than  at  high  ones.  (In  accordance  with  the 
principle  of  Le  Chatelier  (p.  Ill)  is  the  combination  accompanied 
by  the  evolution  or  by  the  absorption  of  heat  ?) 

Speed  of  Chemical  Reactions.  If  hydriodic  acid  is  heated  at 
a  given  temperature,  the  decomposition  seems  to  proceed  more 
and  more  slowly  as  the  reaction  goes  on,  as  will  be  seen  from  the 
following  table : 

RATE  OF  DECOMPOSITION  OF  HYDRIODIC  ACID  AT  374 Ol 


TOTAL  TIME  OF 
HEATING 

FRACTION 
DECOMPOSED 

FRACTION  DECOMPOSED 
IN  1  MINUTE 

360  minutes 

0.0715 

0.00020 

720  minutes 

0.1267 

0.00015 

1080  minutes 

0.1596 

0.00009 

1440  minutes 

0.1715 

0.00003 

A  little  consideration  of  these  results  leads  us  to  the  conclusion 
that  the  decreasing  rate  is  due  to  two  causes  :  first,  because  the 
amount  of  hydriodic  acid  in  the  mixture  is  constantly  decreasing, 
and  second,  because  the  hydrogen  and  iodine  which  result  from 
the  dissociation  are  recombining. 

If  we  begin  with  the  mixture  of  hydrogen  and  iodine,  the  com- 
bination appears  to  be  rapid  at  first,  but  soon  decreases  in  its  rate 
as  the  amounts  of  hydrogen  and  iodine  grow  less  and  as  the 
hydriodic  acid  which  is  formed  begins  to  decompose.  In  order 
to  get  the  real  rate  for  the  decomposition  of  pure  hydriodic  acid 
or  for  the  combination  of  pure  hydrogen  and  iodine,  we  might 
measure  the  rate  for  the  first  infinitesimal  fraction  of  a  minute, 
if  that  were  possible.  For  the  decomposition  at  374°  the  deter- 
mination of  the  composition  of  the  mixture  at  the  end  of  six 
hours  gives  a  rate  for  the  decomposition  of  0.0002,  or  1/5000 

1  Bodenstein,  Z.  physik.  Chem,  29,  295. 


SPEED   OF  REACTIONS  149 

part  of  the  whole  in  one  minute,  and  this  is  comparatively  close 
to  the  rate  of  decomposition  for  pure  hydriodic  acid,  since  at 
the  end  of  the  first  six  hours,  only  about  1/15  of  the  whole  has 
been  decomposed. 

Concentration  and  Speed  of  Reaction.  In  order  to  calculate 
the  true  rate  more  accurately,  it  is  necessary  to  use  the  law  of  the 
relation  between  the  concentration  of  reacting  substances  and  the 
speed  of  reaction,1  which  has  been  based  on  a  careful  study  of 
many  different  reactions  which  take  place  slowly  enough  so  that 
the  rate  can  be  measured.  This  is  that  the  speed  of  any  reaction 
at  a  constant  temperature  (i.e.,  the  part  of  the  whole  which  will 
react  in  unit  time)  is  equal  to  the  product  of  the  concentrations  of 
each  reacting  substance  multiplied  by  a  force  which  is  characteristic 
of  the  given  reaction. 

The  force  which  causes  the  reaction,  and  which  is  given  a 
numerical  value  in  relation  to  the  speed  of  the  reaction  by  this 
law,  is  spoken  of  in  most  textbooks  as  chemical  affinity,  but  it  is 
evidently  complex,  depending  on  the  attraction  between  the 
atoms  which  unite,  the  attraction  between  atoms  which  separate, 
the  temperature,  the  presence  of  catalytic  agents  and  perhaps 
on  many  other  factors.2  The  temperature,  especially,  has  a  large 
effect,  such  that  the  speed  of  a  reaction  is  usually  doubled  for 
an  increase  of  10°. 

For  hydriodic  acid  the  law  may  be  given  the  following  expres- 
sion : 

Let  C-g_t  =  Concentration  of  hydrogen, 
Ci2   =  Concentration  of  iodine, 
CHI  =  Concentration  of  hydriodic  acid, 
FI    =  Force  driving  the  reaction,  H2  +  I2  ^  HI  +  HI,  to 
the  right, 

1  Often  called,  less  correctly,  the  "Law  of  Mass  Action." 

2  In  the  case  of  hydriodic  acid  the  nature  of  the  force  is  wholly 
changed  by  the  action  of  light  and  the  reaction  becomes  unimolec- 
ular: 

HI  =  H  +  I. 

Bodenstein,  Z.  physik.  Chem.  22,  23. 


150 


A  TEXTBOOK  OF  CHEMISTRY 


F2  =  Force  driving  the  reaction  to  the  left, 
Si  =  Speed  of  formation  of  hydriodic  acid, 
82  =  Speed  of  decomposition,  all  at  a  given  temperature 
Then  : 

CW2  x  Cj2  x  FI  =  Si 

CHI  X  CHI  X  FZ  =  Sz l 

By  means  of  these  formulas  it  is  possible  to  calculate  from  the 
results  obtained  by  heating  mixtures  of  hydrogen  and  iodine, 
or  by  heating  hydriodic  acid  for  different  lengths  of  time,  the  rate 
of  combination  or  of  decomposition  for  the  pure  substances  at 
unit  concentration.  The  methods  of  calculation  are  compli- 
cated and  need  not  be  given  here.  The  results  are  as  follows : 

RATE  OF  FORMATION  AND  DECOMPOSITION  OF  HYDRIODIC  ACID 


TEMPER- 
ATURE 

-Si 
FRACTION  OF  WHOLE 
FORMED  IN  ONE  MINUTE 

& 

FRACTION  OF  WHOLE 
DECOMPOSED  IN  ONE 

MINUTE 

Si 

52 

302° 

0.000353 

0.00000326 

1:108 

374° 

0.0140 

0.000221 

1:    63 

427° 

0.172 

0.0031 

1:   55 

It  is  clear  from  this  table  that  the  equilibrium  mixtures  as  given 
on  p.  147  contain  much  more  hydriodic  acid  than  hydrogen 
and  iodine  because  the  combination  takes  place  much  more 
rapidly  than  the  decomposition.  It  is  also  seen  that  the  speed 
of  each  reaction  increases  rapidly  with  the  temperature  and  that 
the  speed  of  decomposition  increases  more  rapidly  than  the  speed 
of  combination.  From  this  it  follows  that  combination  is  ac- 

1  This  depends  on  the  fact  that  the  decomposition  results  from 
the  action  of  two  molecules  of  hydriodic  acid  on  each  other  with  the 
formation  of  molecules  of  hydrogen  and  iodine.  Such  a  reaction 
is  called  bimolecular.  The  action  of  light  seems  to  cause  the  direct 
separation  of  the  atoms  of  hydrogen  and  iodine  from  each  other 
and  the  reaction  in  that  case  is  unimolecular.  See  footnote,  p.  149. 


SPEED  OF  REACTIONS  151 

companied  by  an  evolution  of  heat,  in  accordance  with  the 
principle  of  Le  Chatelier  (see  below  and  also  pp.  Ill  and  201). 

Calculation  of  the  Relative  Speed  of  Two  Reactions  from 
the  Composition  of  an  Equilibrium  Mixture.  If  we  know  the 
nature  of  a  reversible  reaction  and  the  composition  of  the  equilib- 
rium mixture,  it  is  possible  to  calculate  the  relative  speeds  of 
the  opposing  reactions,  at  unit  concentration. 

From  the  formulas  given : 

C\s    /~Y          \.s     If      O 
Hj    s^         Ij       ^          1    —         * 

CHI  X  CHI  X  F 2  =  02 

From  these  equations,  FI  =  Si  and  F*  =  <S2  when  the  concen- 
tration is  1  for  each  of  the  reacting  substances. 

At  374°  the  composition  of  the  equilibrium  mixture  is  very 
near  to : 

80  per  cent  of  HI, 
10  per  cent  of  H2,  and 
10  per  cent  of  I2,  by  volume. 
From  this : 

CHI  =  0.8 
CHa  =  0.1 
Ci,  =  0.1 

Let  S3be  the  speed  of  combination  and  84  the  speed  of  decomposi- 
tion, at  equilibrium.     At  equilibrium  <S3  must  equal  (84. 
Then, 

CH,  X  CIa   X  Fi  =  S, 
0.1  X  0.1  X  Fi  =  S3      • 
CHI  X  CHI  X  F2  =  $4 
0.8  X  0.8  X  F2  =  S4 
Since  <S3  =  S4 

0.1  X  0.1  X  F!  =  0.8  X  0.8  X  F2 

F 

and  — -  =  64.     It  will  be  seen  that  this  result  agrees  closely,  as 

it  should,  with  the  result  obtained  by  the  direct  measurement 
of  the  speeds  of  the  two  reactions  (p.  150).  The  result  may 


152  A  TEXTBOOK  OF  CHEMISTRY 

also  be  checked  by  the  following  calculation,  putting  FI  =  Si 
and  FZ  =  82  and  using  the  values  for  Si  and  $2  given  on  p.  150 : 

CH2  X  Ci2  X  Si  =  S, 
0.1  X  0.1  X  0.0140  =  S3  =  0.000140 

CHI  X  CHI  X  02  =  $4 
0.8  X  0.8  X  0.00022  =  S4  =  0.000140 

This  means,  of  course,  that  after  equilibrium  is  reached  at 
374°,  the  dissociation  and  recombination  still  continue  at  the 
rate  of  about  1/7000  part  of  the  whole 
each  minute. 

Effect  of  Removing  one  of  the  Reacting 
Substances.     Displacement  of  the  Equi- 
librium Point.     If  a  tube  containing  hy- 
™.      _~  driodic  acid  is  heated  in  such  a  way  that 

one  end  of  the  tube  is  kept  cool,  the  iodine 

which  results  from  the  dissociation  will  partly  condense  and  the 
concentration  of  the  iodine,  Ci2,  will  be  diminished  in  the  equation  : 

CH2  X  Cl2  X  Fi  =  S, 

which  gives  the  rate  of  recombination.  Under  these  conditions 
it  is  evident  that  the  decomposition  must  go  much  farther  than 
usual  before  S3  =  84-  In  other  words,  the  removal  of  one  of  the 
constituents  of  a  reversible  reaction  always  displaces  the  equilib- 
rium to  the  side  on  which  the  constituent  removed  appears. 
This  effect  has  been  noticed  '  before  in  the  reaction  between 
iron  and  steam  and  in  that  between  salt  and  sulfuric  acid.  If 
the  end  of  the  tube  could  be  kept  cold  enough  so  that  the  vapor 
pressure  of  iodine  in  it  would  be  reduced  to  zero,  83  would  finally 
become  zero,  and  the  decomposition  of  the  hydriodic  acid  would 
continue  till  it  was  complete. 

Heat  of  Formation  of  Hydriodic  Acid.     The  heats  of  formation 
of  the  compounds  of  the  halogens  with  hydrogen  are  as  follows  : 

H2  +  F2  =  2  X  38,000  calories 
H2  +  C12  =  2  X  21,800  calories 
H2  +  Br2  =  2  X  8300  calories 
Ha  H-  I2  =  2  X  96  calories 


FLUORINE  153 

The  heat  energy  liberated  during  the  combination  grows  less 
with  increasing  atomic  weight  and  becomes  very  small  in  the 
case  of  iodine. 

Bodenstein  calculates  (Z.  physik.  Chem.  29,  313)  the  heat  of 
combination  of  hydrogen  and  iodine  in  gaseous  form  as  follows  • 

At  510°,  2  X  2222  calories 
At  290°,  2  X  943  calories 
At  20°,  2  X  96  calories 

Fluorine,  F,  19.0.  Occurrence.  The  chlorides,  bromides  and 
iodides  of  four  of  the  most  common  metals,  calcium,  magnesium, 
sodium  and  potassium,  are  all  easily  soluble  in  water,  and  these 
three  halogens  are  found  chiefly  associated  with  these  metals,  and 
especially  with  sodium,  in  the  ocean  and  in  brines.  Fluorine,  on 
the  other  hand,  combines  with  calcium  to  form  an  almost  insolu- 
ble compound,  calcium  fluoride,  CaF2,  and  for  this  reason  can 
never  be  found  in  more  than  very  small  amounts  in  natural 
waters,  which  practically  always  contain  calcium.  Fluorine  is 
found  chiefly  as  calcium  fluoride,  CaF2,  in  the  mineral  fluorite. 
Cryolite,  a  double  fluoride  of  aluminium  and  sodium,  NasAlF6 
(or  AlF3.3NaF),  found  in  Greenland,  and  apatite,  a  double  phos- 
phate and  fluoride,  or  chloride,  of  calcium,  Ca6(PO4)3F  or 
Ca5(PO4)3Cl,1  found  in  Canada  and  elsewhere,  are  important 
as  sources  of  aluminium  and  phosphorus  rather  than  as  sources 
of  fluorine. 

Preparation.  When  we  consider  that  the  heat  of  combination 
of  hydrogen  and  fluorine  is  nearly  twice  that  for  the  combination 
of  hydrogen  and  chlorine  (p.  152),  and  remember  that  chlorine 
can  take  a  part  of  the  hydrogen  away  from  oxygen,  we  may  be 
led  to  expect  that  free  fluorine  cannot  exist  in  the  presence  of 
water.  It  was  not  till  this  came  to  be  clearly  understood  that 
Moissan  succeeded  in  obtaining  the  free  element  in  1886.  He  did 
this  by  electrolyzing  a  solution  of  potassium  fluoride,  KF,  in 
anhydrous  hydrofluoric  acid.  He  used  a  U-tube  of  platinum 

1  Note  the  relation  between  this  formula,  the  formula  of  phosphoric 
acid,  H3PO4,  and  the  valences  of  calcium  and  of  fluorine  or  chlorine. 


154  A  TEXTBOOK  OF  CHEMISTRY 

at  first,  but  showed  later  that  a  tube  of  copper  is  only  very 
slightly  attacked  by  the  fluorine,  if  the  temperature  is  kept  at 
—  23°,  or  below,  by  a  freezing  mixture.  In  the  electrolysis, 
fluorine  goes  toward  the  anode  and  is  liberated  there  while 
potassium  and  hydrogen  go  toward  the  cathode,  but  only  hydro- 
gen is  liberated,  because  hydrogen  ions  are  discharged  at  a  much 
lower  potential  than  potassium  ions. 

Properties.  Fluorine  is  a  greenish  yellow  gas,  less  deeply 
colored  than  chlorine.  The  weight  of  a  gram  molecular  volume 
is  38  grams,  showing  that  the  formula  is  F2.  Fluorine  is  the 
most  active  of  the  nonmetallic  elements,  as  is  to  be  expected 
from  its  unique  position  in  the  periodic  system.  It  combines 
directly  and  vigorously  with  nearly  all  elements,  both  metals  and 
nonmetals,  except  with  oxygen.  Many  elements,  as  iodine, 
phosphorus,  arsenic,  carbon  as  charcoal  or  lampblack,  silicon, 
potassium  and  sodium  take  fire  and  burn  in  the  gas,  forming 
fluorides.  Fluorine  will  also  displace  nearly  all  other  non- 
metallic  elements  from  their  compounds.  If  led  into  water, 
it  gives  hydrofluoric  acid  and  oxygen,  rich  in  ozone : 

3H2O  +   3F2   =    O3     +     6HF 

Ozone         Hydrofluoric 
Acid 

Etching  Glass.  Hydrofluoric  Acid  may  be  easily  prepared 
by  warming  a  fluoride  with  concentrated  sulfuric  acid : 

CaF2  +  H2SO4  =  CaSO4  +  2  HF 

Hydrofluoric  acid  is  a  gas  which  may  be  condensed  to  a  liquid 
much  more- easily  than  the  other  halogen  acids.  The  anhydrous 
liquid  boils  at  19.4°.  The  liquid  mixes  with  water  in  all  propor- 
tions, the  concentrated  solution  fuming  in  the  air  in  the  same 
manner  as  concentrated  solutions  of  the  other  halogen  acids. 
The  gaseous  acid  is  very  poisonous  and  the  concentrated  or  anhy- 
drous acid  causes  painful  wounds,  which  are  very  difficult  to 
heal. 

The  most  interesting  property  of  the  acid  is  its  action  on  sili- 


FLUORINE  155 

cates  and  especially  on  glass,  which  is  a  complex  silicate  of 
calcium  and  sodium  or  other  metals.  When  hydrofluoric  acid 
comes  in  contact  with  glass,  the  fluorine  combines  both  with  the 
silicon  and  with  the  metals  of  the  glass  : 

CaSiO3  +  6  HF     =      SiF4     +     CaF2  +  3  H2O 

Calcium  Silicon 

Silicate  Tetrafluoride 

The  reaction  may  be  looked  on  as  a  displacement  of  oxygen 
by  fluorine,  two  atoms  of  fluorine  displacing  one  atom  of  oxygen 
in  accordance  with  the  valences  of  the  two  elements.  Silicon 
tetrafluoride,  SiF4,  is  a  gas  and  escapes.  By  covering  a  glass 
object  with  beeswax,  which  is  not  affected  by  hydrofluoric  acid, 
and  exposing  it  to  the  action  of  the  gas,  after  drawing  lines  or 
figures  through  the  wax  so  as  to  expose  part  of  the  surface  of  the 
glass,  it  is  possible  to  etch  the  exposed  surface  and  obtain  per- 
manent markings  of  any  form  that  is  desired.  Graduation  marks 
on  thermometers,  burettes,  eudiometers,  etc.,  are  made  in  this 
way.  The  best  results  are  obtained  by  exposing  the  glass  to 
the  anhydrous  gas  for  some  hours. 

Commercial  hydrofluoric  acid  is  kept  in  lead  bottles,  which  are 
only  slightly  attacked.  The  pure  acid  must  be  kept  in  platinum 
or  in  bottles  made  of  ceresin,  a  mineral  wax  with  a  higher  melting 
point  than  that  of  paraffin.  The  constant  boiling  solution  boils 
at  120°  and  contains  35  per  cent  of  the  acid. 

Hydrofluoric  acid,  unlike  the  other  halogen  acids,  forms  both 
acid  and  neutral  salts.  Thus  it  forms  with  potassium,  acid 
potassium  fluoride,  KHF2,  as  well  as  the  neutral,  or  normal,  fluo- 
ride, KF.  The  formation  of  these  acid  salts  seems  to  be  closely 
related  to  the  abnormal  density  of  the  gas  and  indicates  that  the 
true  formula  of  the  acid  in  solution  or  at  low  temperatures  is 
probably  H2F2  instead  of  HF.  The  weight  of  a  gram  molecular 
volume  of  gas  varies  from  51.2  grams  at  26°  to  20.6  grams  at  88°. 
At  the  lower  temperature  the  gas  is  evidently  more  complex 
than  H2F2,  for  which  the  gram-molecular-volume  would  weigh 
40  grams. 


156  A  TEXTBOOK   OF  CHEMISTRY 

Metallic  Elements  of  Group  VII.  Manganese  stands  between 
chlorine  and  bromine  in  the  seventh  group  of  the  Periodic  System 
when  the  system  is  given  its  simplest  form  (p.  134).  It  re- 
sembles chlorine  in  the  dioxide,  MnO2,  which  corresponds  to 
chlorine  dioxide,  C1O2,  and  in  permanganic  acid,  HMnQj, 
corresponding  to  perchloric  acid.  But  in  most  of  its  properties 
manganese  is  metallic,  and  it  will  be  considered  further  later 
(p.  533). 

The  Periodic  System  indicates  the  possibility  of  three  or  four 
other  elements  in  the  seventh  group  with  atomic  weights  greater 
than  that  of  bromine,  but  no  such  elements  have  been  dis- 
covered. 

EXERCISES 

1.  Write  the  equations  for  sixteen  reactions  between  the  following 
acids  and  bases,  giving  normal  salts  :  Hydrochloric  acid,  HC1 ;  perchloric 
acid,  HC1O4;  sulfuric  acid,  H2SO4;   phosphoric   acid,  H3PO4;    sodium 
hydroxide,  NaOH;     ferrous    hydroxide,  Fe(OH)2;     ferric   hydroxide, 
Fe(OH)3;    stannic  hydroxide,  Sn(OH)4. 

2.  Write  the  equations  for  the  reactions  between  the  following  salts 
and  sulfuric  acid :   sodium  chloride,  NaCl ;  calcium  chloride,  CaCl2 ; 
sodium  perchlorate,  NaClO4 ;  aluminium  chloride,  A1C13. 

3.  Write  the  equations  for  the  reactions  between  hydrochoric  acid 
and  the  following  oxidizing  agents.     Notice  the  changes  in  valence : 

MnO2     -*•  MnCl2 
KMnO4  ->  KC1  and  MnCl2 
HC10     ->  HC1 
Pb304     ->  PbCl2 
K2Cr207^KClandCrCl3 
KC1O3    ->KC1 

4.  Write  the  equation  for  the  reaction  between  potassium  iodide, 
manganese  dioxide  and  sulfuric  acid,  giving  K2SO4  and  MnSO4. 

5.  Write  the  equation  for  the  reaction  between  calcium  bromide, 
CaBr2,   potassium   permanganate,   KMnO4  and  sulfuric  acid,  giving 
calcium  sulfate  and  the  other  products  to  be  expected. 

These  reactions  are  introduced  here  to  give  the  student  facility  in 
writing  equations  on  the  basis  of  the  valence  of  the  elements.  The 
fundamental  conception  of  valence  is  that  each  atom  has  the  power 
of  holding  directly  in  combination  a  definite,  small  number  of 
other  atoms.  Thus,  when  we  write  the  graphical  formula  H — Cl, 
the  thought  which  it  is  intended  to  convey  is  that  a  hydrogen  or 


THE   HALOGEN  FAMILY  157 

3,  chlorine  atom  holds  directly  to  only  a  single  other  atom  in  the 

/H 

compound,   hydrochloric  acid.     In  water,  H  —  O  —  H    or   (X      ,   the 

H 

oxygen  atom  holds  directly  to  two  other  atoms.     In  some  sense  we 
may  think  that  an  oxygen  atom  has  two  points  of  attachment  for 

/H 


other  atoms.     In  ammonia,  N^-H   or  H  —  N\    ,  in  accordance  with 

\H  XH 

the  same  theory,  each  nitrogen  atom  holds  directly  to  three  hydrogen 
atoms. 

In  the  series  of  oxides  of  the  Periodic  System  the  elements  of  the  zero 
group  do  not  combine  with  other  elements  at  all,  and  these  elements 
are  considered  to  have  a  valence  of  zero.  The  elements  of  the  first 
group  are  univalent,  and  one  bivalent  oxygen  atom  can  hold  two  atoms 
of  these  elements  as  in  Na2O  or  Na  —  O  —  Na.  The  bivalent  atoms  of 
the  second  group  can  hold  bivalent  oxygen  atoms,  atom  for  atom,  as 
in  Mg=O.  In  the  third  group,  where  the  elements  are  trivalent,  if 
we  consider  an  atom  of  such  an  element  as  combined  with  one  atom  of 
oxygen,  one  valence  of  the  first  element  will  remain  unsatisfied,  thus, 

B       .     If  a  second  atom  of  oxygen  is  added,  one  valence  of  this  will 


be  unattached,  B\Q-      On  adding  a  second  atom  of  the  trivalent 
element  and  a  third  atom  of  oxygen,  all  of  the  valences  will  be  balanced 

In  the  next  group  a  quadrivalent  atom  can  balance  two  bi- 

~ 

valent  oxygen  atoms,   Cx'    .     The  same  principles   may   be  easily 


V 
extended  to  the  compounds,  N2O5,  SO3  and  C12O7. 

When  the  oxides  are  those  of  nonmetallic  elements,  they  will,  in  most 
cases,  combine  with  water  to  form  acids.  In  this  case  one  valence  of 
one  oxygen  atom  separates  from  one  nonmetallic  atom,  and  the  hydro- 
gen, H,  of  the  water  attaches  itself  to  the  oxygen,  while  the  hydroxyl, 
OH,  of  the  water  attaches  itself  to  the  nonmetallic  atom : 


H-°X° 

X 


H-0 


158  A  TEXTBOOK  OF  CHEMISTRY 

In  the  reactions  between  acids  and  bases  the  same  principles  of  bal- 
ancing valences  are  to  be  applied,  the  only  difference  being  that  the 
valences  of  the  metal  on  the  one  hand  are  to  be  balanced  against  the 
valences  of  the  acid  groups  on  the  other.  Since  the  hydroxyl  group 
— O — H  is  univalent,  the  number  of  hydroxyl  groups  in  the  base  gives 
the  valence  of  the  metal  of  the  base,  while  the  number  of  replaceable 
hydrogen  atoms  gives  the  valence  of  the  acid  group.  Thus  iron  is  bi- 
valent in  ferrous  hydroxide,  Fe(OH)2,  and  trivalent  in  ferric  hydroxide, 
Fe(OH)3,1  while  the  sulfate  group,  SO4,  of  sulfuric  acid  is  bivalent  and 
the  phosphate  group,  PO4,  of  phosphoric  acid,  H3PO4,  is  trivalent.  By 
representing  the  valences  with  lines,  it  is  a  simple  matter  to  balance 
the  valences  of  a  metal  against  the  valences  of  an  acid  radical  and  so 
determine  the  correct  formula  of  a  salt.  Thus  for  ferric  sulfate  the 
formula  must  be : 


or  for  ferric  phosphate,  Fe=PO4.  A  little  practice  of  this  sort  will 
soon  enable  a  student  to  write  correct  formulas,  such  as  Fe2(SO4)s  or 
FePO4,  without  the  use  of  the  lines  to  indicate  valences.  If  these  prin- 
ciples are  once  understood,  one  needs  to  remember  only  the  formula  of  a 
single  salt  of  any  metal  with  some  well-known  acid  in  order  to  be  able 
to  write  the  formulas  of  the  normal  salts  of  the  metal  with  a  hundred 
or  more  acids  whose  formulas  are  known. 

In  reactions  which  involve  oxidation  and  reduction  it  often  happens 
that  the  valence  of  some  element  changes.  In  a  reduction,  oxygen 
or  some  other  element  is  removed  without  being  replaced,  or  hydrogen 
is  added,  and  to  do  this  hydrogen  must  usually  be  furnished  from 
some  source,  and  the  element  combined  with  this  hydrogen  is  often 
liberated  in  the  free  state. 

Thus  in  the  reaction  between  manganese  dioxide  and  hydrochloric 
acid  quadrivalent  manganese  changes  to  the  bivalent  form.  The  extra 
oxygen  atom  is  balanced  by  hydrogen  from  the  hydrochloric  acid  and  the 
chlorine  of  the  latter  is  liberated.  The  manganese  dioxide  is  reduced, 
the  hydrochloric  acid  is  oxidized: 

O  Cl 

«:£}jMn<:+2H2o+ci2 

H-Cl 

1  For  the  sake  of  simplicity  the  possibility  of  such  doubled 
formulas  as  Fe2(OH)4  and  Fe2(OH)6  is  not  presented  here. 


THE   HALOGEN  FAMILY  159 

In  a  similar  way  if  hydrochloric  acid  acts  on  potassium  permanganate, 
KMnO4,  only  three  chlorine  atoms  are  taken  by  the  potassium  and 
manganese,  and  the  hydrochloric  acid  which  furnishes  these  will  give 
only  three  of  the  eight  hydrogen  atoms  necessary  to  balance  the  four 
oxygen  atoms  of  the  permanganate  molecule.  To  balance  the  remainder 
of  the  oxygen  atoms,  five  more  hydrogen  atoms  will  be  required.  This 
gives  us  the  reaction  : 

KMnO4  +  3  HC1  =  KC1  +  MnCl2  +  4  H2O  +  5  Cl 
+  5HC1 


If  we  wish  to  take  account  of  the  fact  that  free  chlorine  has  the 
formula  C12,  the  equation  must,  of  course,  be  doubled,  giving : 

2  KMn04  +  16  HC1  =  2  KC1  +  2  MnCl2  +  8  H2O  +  5  C12 
Which  substance  is  reduced  and  which  is  oxidized  in  this  reaction  ? 


CHAPTER  XI 
SULFUR,  SELENIUM  AND   TELLURIUM 

THE  nonmetallic  elements  of  Groups  VI  and  VII  of  the  peri- 
odic system  are : 

O    ....     16  F     ....     19 

S     ....     32  Cl  .  .  .  .    35.5 

Se   ....     78  Br  ....     80 

Te  .  .  .  .  127.6  I     ....  127 

Sulfur,  S,  32.0.  Occurrence.  Oxygen  is  found  free  in  nature, 
partly  because  of  its  great  abundance,  forming,  as  it  does,  one 
half  of  that  portion  of  the  earth  which  we  can  examine  directly, 
partly,  probably,  because  of  its  unique  relationship  to  carbon 
and  the  growth  of  plants  (p.  312).  Sulfur  is  also  found  free, 
partly  because  it  is  a  comparatively  abundant  element  and 
partly  because  it  is  easily  liberated  from  hydrogen  sulfide  and 
other  sulfides  by  the  action  of  oxygen  and  some  compounds  of 
oxygen.  Free  sulfur  is  found  in  large  quantities  in  Sicily  and  in 
Louisiana.  Until  about  1903  the  sulfur  mines  of  Sicily  held, 
for  a  long  time,  a  practical  monopoly  of  the  sulfur  markets  of 
the  world,  almost  the  only  competition  coming  from  the  sulfur 
obtained  by  the  Chance  process  (p.  457)  as  a  by-product  in  the 
manufacture  of  sodium  carbonate.  The  sulfur  in  Sicily  is  mixed 
with  other  minerals,  from  which  it  is  separated  by  piling  up  the 
mixture  and  setting  fire  to  the  sulfur  in  such  a  way  that  the 
heat  from  burning  a  part  of  the  sulfur  melts  the  rest  and  the 
latter  runs  out  and  is  collected.  The  process  is,  of  course,  a 
wasteful  one  —  as  a  pound  of  coal  would  give  nearly  as  much 
heat  as  four  pounds  of  sulfur  (p.  27).  The  crude  sulfur  is  re- 
fined by  distillation.  If  the  vapors  are  condensed  in  cold  cham- 
bers, the  sulfur  takes  the  form  of  flowers  of  sulfur,  just  as  the 
freezing  of  water  vapor  gives  snow.  If  the  condensing  room  is 

160 


SULFUR 


161 


hot 
water 


above  the  melting  point  of  sulfur,  the  liquid  sulfur  which  collects 
on  the  bottom  is  run  into  molds  and  forms  the  roll  brimstone  of 
commerce. 

The  extensive  deposits  of  sulfur  in  Louisiana  are  below  a  layer 
of  quicksand,  and  for  a  long  time  after  they  were  discovered  no 
practical  method  of  working  the  deposits  was  known.  The 
difficulty  was  finally  solved  by  a  process  invented  by  Mr.  Frasch 
of  New  York'.  Three  concentric  iron  pipes  are  sunk  to  the  level 
of  the  sulfur,  and  hot  water  under  pressure  is  forced  down  be- 
tween the  two  outer  pipes,  the 
pressure  of  the  water  being  great 
enough  so  that  the  boiling  point 
is  raised  above  the  melting  point 
of  the  sulfur,  114.5°.  The  hot 
water  melts  the  sulfur,  which 
rises  in  the  second  tube,  the  end 
of  which  is  brought  below  the  sur- 
face of  the  melted  sulfur.  To 
bring  the  sulfur  to  the  surface,  air 
is  forced  down  through  the  cen- 
tral tube,  the  sulfur  and  com- 
pressed air  rising  together  between 
the  central  and  second  tubes. 

By  this  process  the  production 
of  sulfur  in  the  United  States  was 
increased  from  3500  tons  in  1900 
to  265,000  tons  in  1911.  The 
world's  production  of  sulfur  in  1909 
was  818,000  tons. 

Sulfur  is  also  found  in  nature 
combined  with  metals  as  metallic 

sulfides  and  with  metals  and  oxygen  as  sulfates.  The  most 
important  sulfides  are  lead  sulfide,  or  galena,  PbS,  zinc  sul- 
fide,  or  sphalerite,  ZnS,  iron  sulfide,  or  pyrite,  FeS2,  and  an 
iron-copper  sulfide,  copper  pyrites,  CuFeS2.  The  most  important 
sulfates  are  calcium  sulfate,  or  gypsum,  CaSO4.2  H2O,  and  barium 


Fig.  51 


162  A  TEXTBOOK  OF  CHEMISTRY 

sulfate,  or  barite,  BaSO4.  Of  these,  only  iron  pyrites  is  used 
primarily  as  a  source  of  sulfur,  for  the  manufacture  of  sulfuric 
acid.  The  other  sulfides  are  used  primarily  for  the  metal  which 
they  contain,  but  sulfuric  acid  is  sometimes  made  from  them  as 
a  by-product. 

Allotropic  Forms  of  Sulfur.  Sulfur  may  exist  in  three  well- 
defined  solid  forms,  in  two  liquid  forms,  which  correspond 
closely  to  two  of  the  solid  forms,  and  in  three  gaseous  forms. 
The  solid  forms  are  : 

1.  Rhombic  Sulfur.     Light  yellow  crystals,  most  often  in  the 
form  of  rhombic  pyramids  (p.  194),  found  in  nature  and  formed 
by  crystallization  from  carbon  disulfide,  in  which  sulfur  is  easily 
soluble.     The  specific  gravity  is    2.06  and  the    melting  point 
114.5°.     This  is  the  most  dense  and  most  stable  form  at  ordinary 
temperatures,  and  the  other  forms  change  to  this  form  more  or 
less  rapidly  at  temperatures  below  96°. 

2.  Monoclinic  Sulfur.      When  melted  sulfur    is  allowed  to 
cool  slowly,  it  crystallizes  in  long,  transparent  needles  of  the 
monoclinic  system  (p.  195).     These  have  a  specific  gravity  of 
1.96  and  melt  at  119°.     This  form  of  sulfur  is  stable  only  at  tem- 
peratures between  96°  and   119°.     At  lower  temperatures  it 
changes  more  or  less  quickly  to  the  rhombic  form.     The  outer 
form  of  the  needles  is  retained,  but  they  become  opaque  and  then 
consist  of  microscopic  crystals  of  the  rhombic  form. 

3.  Amorphous,    Insoluble    Sulfur.     When    sulfur  which    is 
heated  above  160°  is  cooled  quickly  with  care  that  it  does  not 
come  in  contact  with  crystals  of  sulfur,  which  would  cause  a 
rapid  transformation  to  the  crystalline  form,  it  assumes  a  soft, 
plastic  form,  which  hardens  to  a  solid  mass  after  some  hours 
or  days.     If  this  hardened  mass  is  treated  with  carbon  disulfide, 
it  will  be  found  to  be  mostly  insoluble  and  the  insoluble  portion 
is  amorphous,  i.e.  it  has  no  crystalline  structure. 

The  liquid  forms  of  sulfur  are  : 

1.  Mobile  Liquid  Sulfur  (SA).  Between  the  melting  point 
(114.5°  or  119°)  of  either  form  of  sulfur  and  160°  it  forms  a 
mobile,  pale  yellow  liquid. 


SULFUR  163 

2.  Viscous  Liquid  Sulfur  (S^).  When  heated  to  160°,  sulfur 
suddenly  becomes  dark  colored  and  so  viscous  that  a  test  tube 
containing  it  may  be  inverted  without  its  running  out.  If 
heated  to  a  higher  temperature,  the  liquid  becomes  gradually 
somewhat  more  mobile  and  finally  boils  at  444.70.1  The  boiling 
point  is  frequently  used  to  fix  a  point  on  the  scale  of  ther- 
mometers and  pyrometers. 

The  gaseous  forms  of  sulfur  are : 

1.  Sg.     When  sulfur  is  converted  into  a  vapor  at  250°,  under 
low  pressure,  the  weight  of  a  gram  molecular  volume  is  nearly 
256  grams,  indicating  that  there  are  eight  atoms  in  one  molecule 
and  that  the  formula  is  Sg. 

2.  S2.     Even  at  the  boiling  point  (444.7°),  the  weight  of  a  gram 
molecular  volume  of  sulfur  vapor  is  considerably  less  than  256 
grams  and  it  was  formerly  supposed  that  the  formula  at  tempera- 
tures a  little  higher  than  this  was  85.     A  more  careful  study  of 
the  matter  has  demonstrated  that  the  formula  Sg  is  the  true  one 
at  low  temperatures  and  that  the  heavy  molecules  dissociate 
as  the  temperature  rises  until,  at  800°,  the  formula  becomes  S2, 
the  weight  of  a  gram  molecular  volume  at  that  temperature 
being  64  grams.     It  is  still  somewhat  uncertain  whether  the 
larger  molecules   dissociate  directly  into  molecules   of   82   or 
whether  intermediate  molecules  of  84  or  85  are  found.2 

3.  S.     When  sulfur  vapor  is  heated  to  a  very  high  tempera- 
ture (2000°),  it  dissociates  still  further  until  the  gram  molecular 
volume  weighs  only  32  grams  and  the  formula  becomes  8.     We 
may  suppose  that  at  high  temperatures  the  collisions  between 
molecules  become  more  and  more  violent  until,  at  last,  the 
affinity  between  the  atoms  can  no  longer  withstand  the  disrup- 
tive effect  of  the  collisions. 

Properties  and  Uses  of  Sulfur.  Sulfur  burns  readily  in  air 
or  oxygen,  forming  sulfur  dioxide,  SO2,  with  usually  a  small 
amount  of  the  trioxide,  SOs.  The  volume  of  the  sulfur  dioxide 

1  Bulletin  of  the  Bureau  of  Standards,  Vol.  7.  pp.  3  and  129. 

2  See  Premier  and  Schupp,  Z.  physik.  Chem.  68,  144  (1909),  and 
Stafford,  ibid.  77,  66  (1911). 


164  A  TEXTBOOK  OF  CHEMISTRY 

is  almost  the  same  as  the  volume  of  the  oxygen  from  which  it  is 
formed.  (How  does  this  follow  from  Avogadro's  law  and  the 
formulas  of  oxygen  and  sulfur  dioxide  ?)  Sulfur  combines  with 
most  metals  when  heated  with  them,  forming  sulfides.  The 
combination  with  iron  to  ferrous  sulfide,  FeS,  and  with  copper 
to  cuprous  sulfide,  Cu2S,  is  attended  with  considerable  evolution 
of  heat. 

Sulfur  is  burned  to  sulfur  dioxide  for  the  manufacture  of  sul- 
furic  acid,  for  use  in  bleaching  straw  goods,  for  the  "  sulfuring  " 
of  fruit  in  the  process  of  drying,  to  prevent  darkening  and  the 
growth  of  harmful  organisms.  Sulfur  is  also  used  in  the  manu- 
facture of  carbon  disulfide,  of  gunpowder  and  of  india  rubber. 
It  is  used  directly  or  in  a  lime-sulfur  wash  for  application  to 
vines,  fruit  trees,  etc.,  to  prevent  the  growth  of  fungi  or  other 
harmful  organisms. 

Hydrogen  Sulfide,  H2S,  is  found  in  many  natural  waters,  the 
so-called  sulfur  waters.  It  is  formed  by  the  decomposition  of 
organic  matter  containing  sulfur  and  is  one  cause,  though  by  no 
means  the  only  reason,  for  the  disagreeable  odor  of  decayed  eggs 
and  sewage. 

Hydrogen  sulfide  is  formed  when  hydrogen  is  passed  over 
sulfur  heated  to  its  boiling  point,  as  can  be  shown  by  passing  the 
gas,  subsequently,  through  a  solution  of  lead  nitrate,  in  which  it 
will  produce  a  black  precipitate  of  lead  sulfide : 

Pb(N03)2  +  H2S  =  PbS  +  2  HNO3 

Lead  Lead 

Nitrate  Sulfide 

The  reaction  is  reversible  : 

2  H2  -f  S2  ^±  2  H2S 

as  can  be  shown  by  passing  hydrogen  sulfide  through  a  hot  glass 
tube,  in  which  a  ring  of  sulfur  will  be  deposited  beyond  the  point 
that  is  heated. 

Hydrogen  sulfide  is  prepared  in  the  laboratory  by  the  action 
of  hydrochloric  or  sulfuric  acid  on  ferrous  sulfide. 


HYDROGEN   SULFIDE 


165 


FeS  +  2  HC1  =  FeCl2    +  H2S 

Ferrous 
Chloride 


FeS  +  H2SO4  =  FeSO4 

Ferrous 
Sulfate 


H2S 


For  the  preparation  of  the  gas  on  a  small  scale  the  apparatus 
used  for  the  preparation  of  hydrogen  is  suitable.  For  the  use 
of  a  laboratory  the  Parsons  apparatus  (J.  Am.  Chem.  Soc.  25, 
233)  is  better,  because 
the  acid  remains  in  con- 
tact with  the  ferrous 
sulfide  till  the  action  is 
complete.  Hydrochloric 
acid  is  more  satisfactory 
than  sulfuric  acid  for 
such  a  generator,  be- 
cause ferrous  sulfate  is 
less  soluble  than  fer- 
rous chloride  and  some- 
times crystallizes  in  the 
tube  through  which 
the  spent  acid  escapes 
(Fig.  52). 

Hydrogen  sulfide  is  a 
colorless  gas  with  a  very 
disagreeable  odor.  It 
is  quite  poisonous,  if 
breathed  in  more  than 
small  amount.  It  may 
be  condensed  to  a  liquid, 
which  boils  at  —  62°  and 
frozen  to  a  solid,  which 
melts  at  —  85°. 


Fig.  52 


Solution  of  Hydrogen  Sulfide.     Henry's  Law.     One  volume  of 
water  absorbs,  or  dissolves,  4.4  volumes  of  hydrogen  sulfide  at 


166  A  TEXTBOOK  OF  CHEMISTRY 

0°,  3.7  volumes,  at  10°  and  3.1  volumes  at  20°.  The  volume  of 
the  gas  dissolved  is,  between  quite  wide  limits,  independent  of 
the  pressure.  Since  the  weight,  or  amount  of  the  gas  in  a  given 
volume,  is  proportional  to  the  pressure,  it  follows  that  the 
amount  of  the  gas  dissolved  varies  directly  with  the  pressure. 
This  is  known  as  Henry's  Law  (discovered  in  1803).  It  applies 
to  partial  pressures  also.  Thus  if  a  gaseous  mixture  contains 
10  per  cent  by  volume  of  hydrogen  sulfide,  the  amount  dissolved 
from  such  a  mixture  at  20°  will  be  only  0.31  volume.  One  hun- 
dred cubic  centimeters  of  water  in  contact  with  pure  oxygen  dis- 
solve 4.9  cc.  of  the  gas  at  0° ;  in  contact  with  nitrogen  100  cc. 
dissolve  2.35  cc.  of  nitrogen.  One  hundred  cubic  centimeters 
of  water  in  contact  with  air  will  contain,  therefore,  4.9  X  0.21 
=  1.04  cc.  of  oxygen  and  2.35  X  0.78  =  1.83  cc.  of  nitrogen. 
The  law  does  not  hold  for  gases  which  are  very  easily  soluble 
in  water,  such  as  hydrochloric  acid  or  ammonia. 

In  accordance  with  Henry's  law,  water  containing  hydrogen 
sulfide  loses  the  gas  rapidly  on  exposure  to  the  air,  in  which  the 
partial  pressure  of  the  gas  is,  of  course,  zero.  In  addition  to 
this  the  oxygen  absorbed  by  the  water  reacts  with  the  hydrogen 
sulfide,  liberating  sulfur : 

2  H2S  +  O2  =  2  H2O  +  2  S 

The  action  is  similar  to  the  liberation  of  chlorine  from  chlorides 
by  fluorine,  but  is  far  less  rapid. 

Sulfides.  Groups  of  Analytical  Chemistry.  When  hydrogen 
sulfide  is  passed  into  a  neutral  or  slightly  acid  solution  containing 
salts  of  certain  metals,  such  as  arsenic,  mercury  and  lead,  the 
metal  is  precipitated  as  a  sulfide  because  the  sulfides  of  these 
metals  are  extremely  insoluble : 

2  AsCl3  +  3  H2S  =  As2S3  +  6  HCl 

HgCl2  +  H2S  =  HgS  +  2  HCl 
Pb(N03)2  +  H2S  =  PbS  +  2  HN03 

If  hydrogen  sulfide  is  passed  into  an  alkaline  solution  contain- 
ing the  salts  of  some  other  metals,  such  as  iron,  zinc  and  manga- 
nese, which  are  not  precipitated  from  acid  solutions,  these  metals, 


STRENGTH   OF  ACIDS  167 

whose  sulfides  are  also  very  insoluble,  but  more  soluble  than 
those  of  the  metals  of  the  first  group,  are  precipitated  also  as 
sulfides. 

FeSO4  +  H2S  +  2  NaOH  (or  Na2S)  =  FeS  +  Na2SO4  +  2  H2O 
ZnS04  +  Na2S  =  ZnS  +  Na2SO4 
MnS04  +  Na2S  =  MnS  +  Na2SO4 

A  part,  but  not  all,  of  the  metals  of  the  first  group  are  precipi- 
tated from  alkaline  as  well  as  from  acid  solutions.  The  reason 
for  the  exceptions  need  not  be  discussed  here.  (See  p.  261.) 

A  third  class  of  metals  form  salts  which  are  not  precipitated 
from  acid,  neutral  or  alkaline  solutions. 

The  conduct  of  solutions  of  metals  toward  hydrogen  sulfide, 
as  just  outlined,  is  the  basis  for  the  separation  of  metals  into  three 
fundamental  groups  for  the  purposes  of  analytical  chemistry. 

Hydrosulfuric  Acid.  Strength  of  Acids.  A  solution  of  hydro- 
gen sulfide  in  water  will  redden  blue  litmus  paper  and  will  neutral- 
ize a  solution  of  sodium  hydroxide,  or  in  other  words  a  certain 
amount  of  a  solution  of  a  base  must  be  added  before  the  hy- 
droxide will  turn  the  litmus  blue.  These  are  the  properties  of 
an  acid,  and  hydrogen  sulfide  is  sometimes  very  properly  called 
hydrosulfuric  acid,  just  as  hydrogen  chloride  is  called  hydrochloric 
acid.  We  have  seen  that  a  solution  containing  a  milligram 
molecule  of  hydrochloric  acid  in  10  cc.  of  water  freezes  at 
—  0.355°,  while  a  solution  of  alcohol  containing  a  milligram 
molecule  in  10  cc.  freezes  at  —  0.184,  and  the  difference  was 
explained  by  supposing  that  the  hydrochloric  acid  separates 
largely  into  hydrogen  (H+)  and  chloride  (Cl~)  ions.  A  solution 
of  hydrogen  sulfide  which  contains  one  milligram  molecule  in 
10  cc.  freezes  at  —  0.196°. 

36.5  mg.  HC1  in  10  cc.  of  H2O  freezes  at  -  0.355 
46  mg.  C2H6O  in  10  cc.  of  H2O  freezes  at  -  0.184 
34  mg.  H2S  in  10  cc.  of  H2O  freezes  at  -  0.196 

This  indicates  that  such  a  solution  of  hydrogen  sulfide  con- 
tains comparatively  few  hydrogen  ions.  This  conclusion  is 


168  A  TEXTBOOK  OF  CHEMISTRY 

confirmed  by  the  electrical  conductivity  of  the  solution.  The 
solution  of  hydrochloric  acid  referred  to  is  a  very  much  better 
conductor  (nearly  2000  times)  of  electricity  than  the  solution  of 
hydrogen  sulfide. 

During  a  long  period  in  the  history  of  chemistry  acids  were 
spoken  of  as  strong  or  weak  according  to  whether  they  could 
expel  other  acids  from  their  salts  or  not.  Thus  sulfuric  acid  was 
thought  to  be  stronger  than  hydrochloric  or  nitric  acid  because 
it  would  expel  these  acids  from  salt  or  saltpeter.  We  have  seen 
that  such  a  view  can  no  longer  be  held  (p.  119)  and  that  all  such 
reactions  are  reversible.  There  is  another  sense,  however,  in 
which  some  acids  are  strong  while  others  are  weak,  and  the  basis 
for  a  true  distinction  of  this  kind  has  just  been  indicated.  A 
strong  acid  is  one  which  separates  largely  into  hydrogen  ions  and 
negative  ions  in  an  aqueous  solution.  A  weak  acid  is  one  that 
separates  to  only  a  comparatively  small  degree  into  hydrogen 
ions  and  negative  ions  in  solution.  In  this  sense  hydrochloric 
acid  is  one  of  the  strongest  of  the  acids,  sulfuric  acid  is  weaker 
but  still  a  very  strong  acid,  hydrofluoric  acid  is  much  weaker, 
acetic  acid  is  still  weaker  and  hydrosulfuric  acid,  H2S,  is  very 
weak  indeed. 

Acids  like  hydrosulfuric  acid  which  contain  two  hydrogen 
atoms  may  ionize  in  either  of  two  ways  : 


or 

In  the  case  of  weak  acids  the  ionization  probably  takes  place 
almost  exclusively  in  the  first  form.  It  is  worthy  of  notice  that 
the  halogen  acids  (hydrochloric  acid,  etc.),  which  contain  only 
one  hydrogen  atom  in  the  molecule,  ionize  very  completely  in 
moderately  dilute  solutions,  while  hydrogen  sulfide,  with  its  two 
hydrogen  atoms,  ionizes  to  only  a  slight  extent.  If  oxygen  is 
added,  however,  as  in  sulfuric  acid,  H^SCX,  the  ionization  be- 
comes large,  though  it  does  not  equal  that  of  hydrochloric  acid. 
Application  of  the  Idea  of  Strength  of  Acids  to  explain  the 


STRENGTH  OF  ACIDS  169 

Conduct  of  Sulfides.1  Practically  all  of  the  ordinary  reactions 
in  aqueous  solutions  are  reversible.  A  reversible  reaction  leads 
to  a  stable  condition  only  when  the  reaction  has  reached  a  point 
where  it  proceeds  just  as  fast  in  one  direction  as  in  the  other. 
In  the  reversible  reactions  : 

Pb++      +  2  NOr  +  H+  +  HS-  ^±  Pb+        +  2  NO3~  +  H+  and 

XSH 


b+ 
X 


+HS-      ^±  PbS  +  H2S2 

SH 


the  reactions  will  proceed  toward  the  right  as  long  as  lead  sul- 
fide  (PbS)  is  formed  and  separates  from  the  solution.  The 
equilibrium  finally  reached  must  depend  upon  whether  there 
are  enough  lead  ions  (Pb++),  lead  hydrosulfide  (Pb+  )  ions  and 

XSH 

hydrosulfide  ions  (HS~)  in  a  given  volume  of  the  solution  to  form 
more  lead  sulfide  than  can  remain  in  solution.     As  lead  sulfide 
is  very  insoluble,  only  a  very  few  hydrosulfide  ions  can  remain 
in  a  solution  containing  lead  ions. 
The  ionization  of  hydrogen  sulfide  : 


takes  place  to  a  very  limited  extent  even  in  pure  water,  the 
equilibrium  in  the  reaction  of  ionization  being  very  far  to  the 
left.  If  we  add  to  a  solution  of  hydrogen  sulfide  a  strong  acid, 
as  hydrochloric  acid,  which  gives  a  large  number  of  hydrogen 
ions,  the  hydrosulfide  ion  (HS~)  will  meet  hydrogen  ions  more 
frequently  than  before  and  will  combine  with  them  to  form 
hydrogen  sulfide.  This  must  shift  the  equilibrium  to  the  left 
and  cause  an  increase  in  the  unionized  hydrogen  sulfide  and  a 

1  The  student  should  read  this  paragraph,  but  it  may  be  well  to 
leave  its  careful  study  till  review  or  a  later  period.    (See  pp.  379-386.) 

2  This  may  involve  the  further  ionization, 

Pb+          ^±  Pb+        +  H+ 

XSH  \S- 

but  the  ion  Pb+       if  capable  of  existence  at  all,  would  immediately 


become  PbS.  Most  authors  are  accustomed  to  write  the  reaction  : 
Pb++  +  2  NOr  +  2  H+  +  S—  =  PbS  +  2  NO8~  +  2  H+,  but  the 
form  given  above  seems  more  probable. 


170  A  TEXTBOOK  OF  CHEMISTRY 

decrease  in  the  number  of  hydrosulfide  ions1  (HS~).  The 
presence  of  a  moderate  amount  of  hydrochloric  acid  in  a  solu- 
tion containing  a  lead  salt  will,  therefore,  so  far  decrease  the  con- 
centration of  the  hydrosulfide  ions  that  lead  sulfide  can  no  longer 
be  precipitated. 

It  will  be  seen  from  what  has  just  been  said  that  the  distinction 
between  the  first  and  second  classes  of  metals  (p.  166)  in  qualita- 
tive analysis  depends  on  our  definition  of  a  <k  slightly  acid  " 
solution.  Such  metals  as  lead,  cadmium  and  zinc  might  belong 
to  the  first  class  or  the  second  according  to  the  concentration 
of  the  hydrogen  ions  present. 

The  addition  of  an  alkali,  as  sodium  hydroxide  (NaOH),  to  a 
solution  of  hydrogen  sulfide  has  an  effect  opposite  to  the  addition 
of  an  acid.  The  base  gives  hydroxide  (OH~)  ions,  which  combine 
with  the  hydrogen  ions  to  form  water.  This  displaces  the 
equilibrium  for  the  ionization  of  hydrogen  sulfide  in  the  oppo- 
site direction  and  results  in  a  large  increase  in  the  number  of 
hydrosulfide  ions  (HS~~).  Under  these  conditions  the  sulfides 
of  iron,  zinc  and  some  other  metals,  which  are  too  soluble  to 
form  at  all  in  acid  solutions,  will  form  and  be  precipitated. 

When  hydrogen  sulfide  is  passed  into  a  solution  containing 
a  hydroxide  of  a  metal  of  the  third  class,  a  hydrosulfide,  which 
ionizes  to  a  large  extent  in  dilute  solutions,  is  formed  : 

Na+  +  OH-  +  H+  +  HS-  =  Na+  +  HS~  +  H2O 

If  a  second,  equal  amount  of  sodium  hydroxide  is  added  and 
the  solution  is  evaporated  to  dryness,  sodium  sulfide,  Na2S, 
may  be  obtained  : 

Na+  +  HS-  +  Na+  +  OH~  ^±  Na2S  +  H2O 

1  This  may  be  stated  mathematically  as  follows : 

Ce+  X  CHS-  X  Fi  =  Si 

CH?s  XFi  '  =  S2 

Since  Si  =  S2  at  equilibrium  and  S2  must  be  constant  for  a  given 
quantity  of  hydrogen  sulfide  and  Fi  is  also  constant,  the  product 
of  CH+  X  CHS~~  must  be  constant  for  any  given  concentration 
of  hydrogen  sulfide.  Any  increase  in  the  number  of  hydrogen  ions 
must,  therefore,  be  accompanied  by  a  corresponding  decrease  in  the 
number  of  hydrosulfide  ions. 


{REDUCTION  BY  HYDROGEN  SULFIDE  171 

If  the  sodium  sulfide  is  dissolved  in  water,  the  ionization  of 
water  approaches  so  near  in  degree  to  that  of  hydrogen  sulfide 
that  the  sodium  sulfide  is  largely  hydrolyzed : 

Na2S  +  H+  +  OH-  =  2  Na+  +  HS~  +  OH~ 

The  presence  of  hydroxide  ions  in  such  a  solution  is  indicated 
by  the  alkaline  reaction  of  the  solution,  as  shown  by  litmus  or 
other  test  papers. 

Hydrogen  Sulfide  as  a  Reducing  Agent.  Hydrogen  sulfide 
readily  gives  up  its  hydrogen  to  chlorine,  bromine  or  iodine. 
It  also  gives  up  hydrogen  to  a  great  variety  of  compounds,  re- 
ducing them.  The  following  are  typical  illustrations  : 

H2S  +  I2  =  2  HI  +  S 

This  reaction  furnishes  an  excellent  method  of  preparing  a  solu- 
tion of  hydriodic  acid,  by  suspending  iodine  in  water  and  passing 
hydrogen  sulfide  into  the  mixture. 

K2Cr207  +  8  HC1  +  3  H2S  -  2  KC1  +  2  CrCl3  +  7  H2O  +  3  S 

Potassium  Chromic 

Dichromate  Chloride 

In  writing  the  equation  for  this  reaction,  notice  that  the  formulas 
of  the  chlorides  determine  the  number  of  molecules  of  hydro- 
chloric acid  required.  Comparing  the  number  of  molecules  of 
hydrochloric  acid  with  the  number  of  atoms  of  oxygen  in  the 
dichromate  it  is  seen  that  after  water  has  been  formed  from  the 
hydrogen  of  the  hydrochloric  acid  three  atoms  of  oxygen  remain. 
These  will  oxidize  the  hydrogen  of  three  molecules  of  hydrogen 
sulfide. 

Fe2(S04)3  +  H2S  =  2  FeSO4  +  H2SO4  +  S 

Ferric  Ferrous  . 

Sulfate  Sulfate 

Here  the  hydrogen  of  the  hydrogen  sulfide  takes  the  sulfate 
radical  (SO4)  from  the  ferric  sulfate,  and  the  iron  is  reduced  from 
the  ferric  to  the  ferrous  state. 

Another  method  of  writing  such  equations,  which  is  preferred 
by  some  teachers,  is  based  on  the  principle  of  positive  and  nega- 
tive valences.  According  to  this  principle  : 


172  A   TEXTBOOK  OF  CHEMISTRY 

1.  The  algebraic  sum  of  the  valences  of  any  compound  is  zero. 
The  valence  of  a  free  element  is  also  zero. 

2.  Oxygen  in  compounds  has  a  negative  valence  of  2. 

3.  Hydrogen  in  compounds  has  a  positive  valence  of  1. 

4.  When  the  valence  of  one  element  changes,  the  valence  of 
some  other  element  or  elements  must  change  by  the  same  amount 
in  the  opposite  direction. 

In  applying  these  principles  to  the  present  case,  the  equation 
is  first  written  in  the  following  form : 

K2Cr2O7  +  HC1  +  H2S  -+  2KC1  +  2  CrCl3  +  S  +  H2O 

On  inspection  it  is  seen  that  the  sum  of  the  valences  of  the  two 
potassium  and  two  chromium  atoms  on  the  left  is  +  14,  while 
on  the  right  the  sum  of  the  valences  of  the  same  four  atoms  is 
only  +  8,  a  loss  of  6  positive  valences.  To  balance  this  the 
valence  of  the  sulfur  atom  changes  from  —  2  in  hydrogen  sul- 
fide,  H2S,  to  0  in  free  sulfur,  S.  It  is  obvious,  at  once,  that 
to  balance  the  changes  in  the  chromium  we  must  have  three 
molecules  of  hydrogen  sulfide.  To  furnish  the  chlorine  for  the 
chlorides  there  must  be  8  molecules  of  hydrochloric  acid,  HC1. 
The  equation  becomes,  therefore  : 

K2Cr2O7  +  8  HC1  +  3  H2S  =  2  KC1  +  2  CrCl3  +  7  H2O+3  S 
Similar  reactions  occur  between  hydrogen  sulfide  and  chlorine, 
hydrogen  sulfide  with  sulfuric  acid  and  potassium  permanganate, 
KMnO4,  or  hydrogen  sulfide  and  ferric  chloride,  FeCls.  The 
reaction  between  lead  sulfide,  PbS,  and  nitric  acid,  HNO3, 
giving  lead  nitrate,  Pb(NO3)2,  nitric  oxide,  NO,  sulfur  and 
water,  is  also  closely  related  to  these,  the  positive,  bivalent  lead 
atom  taking  the  place  of  the  two  positive  hydrogen  atoms  in 
hydrogen  sulfide.  The  student  is  advised  to  write  the  equa- 
tions for  these  reactions  by  use  of  both  of  the  methods  suggested 
above. 

Sulfur  Dioxide  is  formed  when  sulfur  is  burned  in  the  air, 
also  when  iron  pyrites,  FeS2,  is  burned,  the  latter  method  of 
preparation  being  used  largely  in  the  manufacture  of  sulfuric 
acid. 


SULFUR  DIOXIDE  173 

Sulfur  dioxide  may  be  prepared  in  the  laboratory  by  the  reduc- 
tion of  concentrated  sulfuric  acid  with  copper  or  other  sub- 
stances. If  copper  is  used,  copper  sulfate  is  formed.  The 
equation  may  be  written  as  follows  : 


Cu  +  H2SO4  =  [CuO]   +  SO2  +  H2O 
[CuO]  +  H2SO4  =  CuSO4  +  H20 
Combining,      Cu  +  2  H2SO4  =  CuSO4  +  SO2  +  2  H2O 

The  first  two  equations  are  written  as  an  aid  to  the  writing  of 
the  last.  The  [CuO]  is  placed  in  brackets  to  indicate  that  it  is 
not  a  final  product  of  the  reaction.  It  may  or  may  not  be  formed 
as  an  intermediate  product.  Another  method  of  writing  would 
be  to  represent  hydrogen  [2  H]  as  an  intermediate  product. 

The  most  convenient  laboratory  method  for  the  preparation 
of  sulfur  dioxide  is  to  drop  concentrated  sulfuric  acid  into  a  40 
per  cent  solution  of  acid  sodium  sulfite,  NaHSOa  : 

NaHSO3  +  H2SO4  ^±  NaHSO4  +  H2SO3 

Sulfurous 
Acid 


Sulfurous  acid,  H2SOs,  decomposes  very  easily  into  sulfur 
dioxide  and  water,  the  sulfur  dioxide  escaping  as  a  gas.  This 
has  the  same  effect  on  the  equilibrium  of  the  first  reaction  as  if 
the  sulfurous  acid  itself  were  volatile  and  escaped  from  the  mix- 
ture. 

Sulfur  dioxide  is  a  colorless  gas  with  a  suffocating  odor,  famil- 
iar in  the  burning  of  sulfur-tipped  matches.  It  may  be  con- 
densed to  a  liquid  in  a  tube  surrounded  with"  a  freezing  mixture 
and  boils  at  —  10°.  It  freezes  at  a  very  low  temperature  and 
melts  at  -  73°. 

Sulfur  dioxide  is  used  to  bleach  straw,  wool  and  silk.  The 
latter,  especially,  are  injured  by  the  action  of  chlorine,  so  that  it 
cannot  be  used.  The  sulfur  dioxide  seems  to  combine  with  the 
coloring  matter  to  form  colorless  compounds,  or,  in  some  cases, 
to  reduce  the  colored  compound  to  a  colorless  one.  Exposure 


174  A  TEXTBOOK  OF  CHEMISTRY 

to  the  air  and  light  frequently  restores  the  color,  as  in  the  case 
of  straw  hats. 

Sulfur  dioxide  and  sulfites  are  powerful  germicides.  Its  use 
as  a  disinfectant,  however,  has  been  almost  entirely  replaced  by 
formaldehyde,  which  is  even  more  effective  and  does  not  injure 
fabrics  or  metallic  articles,  as  sulfur  dioxide  does.  The  injury  to 
fabrics  may  be  either  through  its  bleaching  effect  or  because 
it  is  slowly  oxidized  by  the  action  of  air  and  moisture  to  sul- 
furic  acid,  which  is  corrosive.  Sulfur  dioxide  is  still  used  exten- 
sively in  "  sulfuring  "  fruit  to  destroy  the  organisms  which  cause 
darkening  and  injury  during  the  drying. 

Sulfurous  Acid.  At  ordinary  temperatures  water  dissolves 
about  50  times  its  volume  of  sulfur  dioxide.  The  solution  red- 
dens litmus  and  neutralizes  bases,  showing  that  the  sulfur  dioxide 
combines  with  the  water  and  forms  an  acid.  The  sodium  salt 
obtained  by  neutralizing  the  acid  is  sodium  sulfite,  Na2SO3, 
and  from  the  formula  of  this  and  other  salts  it  is  assumed  that 
the  formula  of  the  sulfurous  acid  in  such  a  solution  is  H2SOa. 
The  structure  of  the  acid  is  probably 

(k       /H 


CT     \)H 


It  is  very  unstable,  one  hydroxyl  group  and  one  hydrogen  atom 
separating  very  easily  from  the  molecule.  The  solution  smells 
strongly  of  sulfur  dioxide,  and  all  of  the  gas  can  be  expelled  by 
boiling  the  solution. 

Sulfurous  acid  is  a  comparatively  weak  acid.  In  a  solution 
containing  0.05  gram  molecule  (-£$  mol)  about  20  per  cent  of 
its  hydrogen  is  ionized,1  while  in  a  corresponding  solution  of 
sulfuric  acid  60  per  cent  of  the  hydrogen  is  ionized.  Sulfurous 
acid  is  a  powerful  reducing  agent.  It  is  oxidized  to  sulfuric  acid 

1  This  is  on  the  supposition  that  all  of  the  sulfur  dioxide  has 
combined  with  water  to  form  sulfurous  acid.  It  is  probable  that 
some  of  the  sulfur  dioxide  exists  as  such  in  the  solution  and  that  the 
ionization  of  the  sulfurous  acid  really  present  is  considerably  greater 
than  appears  from  these  figures. 


SULFUR  TRIOXIDE  175 

by  potassium  permanganate,  KMnO4,  potassium  dichromate, 
K2Cr2C>7,  chlorine,  bromine  or  ferric  salts. 

Sulfites.  Sulfurous  acid  forms  both  acid  and  normal  salts, 
the  salts  of  sodium  being  acid  sodium  sulfite,  NaHSO3,  and  nor- 
mal sodium  sulfite,  Na2SO3.  The  calcium  salts  are  CaH2(SO3)2 
and  CaSO3.  These  salts  are  prepared,  commercially,  by  burn- 
ing sulfur  in  air  and  passing  the  mixture  of  sulfur  dioxide 
and  nitrogen  through  a  solution  of  sodium  carbonate  or  sodium 
hydroxide  for  the  sodium  salts,  or  through  milk  of  lime  (Ca(OH)2) 
for  the  calcium  salts.  The  acid  sodium  salt  has  been  used  as  an 
addition  to  wine  or  cider  to  stop  fermentation.  The  acid  cal- 
cium salt  is  used  in  the  purification  of  wood  pulp  for  the  manu- 
facture of  paper. 

Sulfur  Trioxide.  Some  heat  is  evolved  when  sulfur  dioxide 
combines  with  oxygen  to  form  the  trioxide,  SO3,  but  the  speed 
of  the  reaction  between  the  two  is  too  slow  to  be  measured  at 
ordinary  temperatures.  At  temperatures  where  the  speed  of 
the  reaction  of  combination  becomes  sufficiently  rapid  to  become 
a  practicable  method  of  preparation,  the  dissociation  of  sulfur 
trioxide  into  sulfur  dioxide  and  oxygen  becomes  very  large  and 
renders  this  method  of  preparation  from  the  substances  alone 
impracticable.  The  reversible  reaction  : 

2  SO2  +  O2  ^±  2  SO3 

has  its  point  of  equilibrium  shifted  toward  the  left  as  the  tem- 
perature rises,  in  accordance  with  the  principle  of  van't  Hoff- 
LeChatelier  (p.  111). 

As  early  as  1831  it  was  discovered  that  the  reaction  is  greatly 
accelerated  by  the  presence  of  platinum,  but  it  was  nearly 
70  years  before  the  details  for  the  application  of  this  principle 
were  so  far  worked  out  as  to  render  the  manufacture  on  a 
large  scale  possible,1  so  long  does  it  often  take  to  convert  a 
scientific  discovery  into  commercial  success.  The  chief  diffi- 
culties to  be  overcome  were,  first,  that  arsenic  and  other  sub- 

1  For  very  interesting  historical  details  see  Knietsch,  Ber.  34, 
4069  (1901). 


176 


A  TEXTBOOK  OF   CHEMISTRY 


stances  in  the  gases  obtained  by  roasting  pyrites  "  poison  " 
the  platinum  and  render  it  ineffective  for  the  catalysis,  and 
second,  that  the  platinum  catalyzes  the  dissociation  of  sulfur 
trioxide  as  well  as  the  combination  of  sulfur  dioxide  and  oxygen, 
and  the  temperature  for  rapid  combination  lies  very  close  to  a 
temperature  at  which  the  dissociation  is  large  and  so  the  com- 
bination becomes  incomplete.  These  difficulties  have  been 
overcome  by  a  careful  purification  of  the  sulfur  dioxide  as  it 
comes  from  the  pyrites  burners  and  by  a  careful  regulation 
of  the  temperature  as  the  gases  pass  over  the  "  contact  mass." 
The  platinum  of  the  "  contact  mass"  is  disseminated  in  a  very 
finely  divided  condition  over  asbestos  or  some  other  material 
which  gives  it  a  large  surface  in  proportion  to  its  weight. 

In  the  laboratory,  on  a  small  scale,  sulfur  trioxide  can  be  read- 
ily prepared  by  passing  dry  sulfur  dioxide  and  oxygen  through  a 
gently  warmed  tube  containing  platinized  asbestos  (Fig.  53). 


Fig.  53 

It  may  be  obtained  still  more  easily  by  warming  "  fuming  " 
sulfuric  acid,  which  is  a  mixture  of  sulfuric  acid,  H2SO4,  pyro- 
sulfuric  acid,  H2S2O7,  and  sulfur  trioxide. 

Sulfur  trioxide  is  a  clear,  volatile  liquid  which  solidifies  at  a 
low  temperature.  It  melts  at  14.8°  and  boils  at  46°.  In  the 
presence  of  a  trace  of  moisture  a  little  sulfuric  acid,  H2SO4,  or 
pyrosulfuric  acid,  H2S2O7,  is  formed  and  this  acts  as  a  catalytic 
agent  causing  sulfur  trioxide  to  polymerize,  forming  the  com- 
pound S2Oe,  which  crystallizes  in  white,  asbestos-like  needles. 
As  it  is  extremely  difficult  to  exclude  moisture  completely,  this 
polymeric  form  is  usually  obtained  instead  of  the  true  trioxide. 
On  warming  it  gives  a  vapor,  which  consists  of  the  true  trioxide. 


SULFURIC  ACID  177 

Sulfur  trioxide  hisses  like  a  hot  iron  when  thrown  into  water, 
owing  to  the  heat  generated  when  it  combines  with  water  to 
form  sulfuric  acid.  It  fumes  strongly  in  the  air,  forming  minute 
drops  of  sulfuric  acid,  which  settle  only  very  slowly  and  are  not 
readily  absorbed  by  water.  Curiously  enough  these  minute 
drops  are  easily  absorbed  by  concentrated  sulfuric  acid  and  this 
is  used  for  the  purpose  in  the  manufacture  of  sulfuric  acid  by 
the  contact  process. 

Sulfuric  Acid.  The  contact  process  for  the  preparation  of 
sulfur  trioxide  has,  thus  far,  been  used  almost  exclusively  for  the 
manufacture  of  a  very  concentrated  or  a  "  fuming  "  sulfuric  acid. 

It  has  been  pointed  out  that  the  direct  combination  of  sulfur 
dioxide  and  oxygen  is  too  slow  to  be  commercially  possible  as  a 
method  of  manufacture,  and  that  platinum  is  used  to  catalyze, 
or  hasten,  the  reaction.  Another  catalytic  agent,  not  so  sensitive 
to  impurities  in  the  gases,  or  to  temperature  changes,  and  which 
acts  rapidly  at  ordinary  temperatures,  has  been  used  for  a  long 
time  in  what  is  called  the  "  chamber  process  "  for  the  manufac- 
ture of  sulfuric  acid.  In  this  process  large  chambers  lined  with 
sheet  lead,  which  is  only  slightly  attacked  by  dilute  sulfuric  acid, 
are  employed.  Into  these  chambers  are  introduced : 

1 .  Sulfur  dioxide  from  burning  sulfur  or  iron  pyrites : 

2  FeS2  +11O  =  Fe2O3  +  4  SO2 

2.  Nitric  acid  from  Chile  saltpeter,  NaNO3,  and  sulfuric  acid  : 

NaNO3  +  H2SO4  =  NaHSO4  +  HNQ3 

3.  Air,  to  furnish  oxygen. 

4.  Water  as  steam  or  spray. 

The  first  reaction  consists  in  the  oxidation  of  the  sulfur  dioxide 
to  sulfuric  acid  by  the  nitric  acid  : 

3  SO2  +  2  HNO3  +  2  H2O  =  3  H2SO4  +  2  NO 1 

1  This  equation  should  not  be  learned  by  rote,  but  should  be  written 
on  the  following  considerations : 

1.  When  nitric  acid  is  reduced  to  nitric  oxide,  two  molecules 
give  3  atoms  of  available  oxygen. 

2.  Each  atom  of  oxygen  will  oxidize  one  molecule  of  sulfur 
dioxide. 


178  A  TEXTBOOK  OF  CHEMISTRY 

If  it  were  necessary  to  stop  here  and  the  nitric  oxide  were  lost, 
sulfuric  acid  would  be,  comparatively,  an  expensive  substance  on 
account  of  the  limited  supply  and  relatively  high  price  of  sodium 
nitrate.  But  nitric  oxide  combines  almost  instantly  with  the 
oxygen  of  the  air  to  form  nitrogen  dioxide,  NC>2  : 

2  NO  +  O2  =  2  NO2 

Nitrogen  dioxide,  in  turn,  can  oxidize  a  new  quantity  of  sulfur 
dioxide  : 

S02  +  NO2  +  H2O  =  H2SO4  +  NO 

It  is  pretty  certain  that  the  mechanism  of  the  reaction  is  more 
complicated  than  is  indicated  by  these  equations,  but  the  equa- 
tions given  indicate  clearly  the  fundamental  facts  on  which  the 
action  depends.  These  are  :  first,  that  nitric  oxide  combines, 
practically  instantaneously,  with  oxygen  ;  second  that  nitrogen 
dioxide  can,  directly  or  indirectly  and  very  quickly,  give  its 
oxygen  to  the  sulfur  dioxide  and  water,  converting  these  to 
sulfuric  acid.  Commercially,  the  whole  process  depends  on  the 
speed  with  which  these  actions  occur. 

The  theory  of  the  lead  chamber  process  which  has  received 
most  acceptance  is  that  of  Lunge,  who  supposes  the  process  to 
consist  in  the  formation  and  decomposition  of  nitrosyl  sulfuric 

add:  NO 

OH 

a  mixture  of  nitric  oxide,    NO,    and   nitrogen    dioxide,   NO2 
(equivalent  to  nitrous   anhydride,  N2Os),  being  the  effective 


2  SO2  +  N2O3  +  O2  +  H2O  =  2  S02  < 

XOH 

X)—  NO  /OH 

2  SO2<  +  H2O  =  2  S02<         +  N2O3 

XOH  XOH 

Nitrosyl  sulfuric  acid  is  a  definite,   crystalline  compound, 
which  is  formed  in  the  chambers  when  the  supply  of  water  is 


SULFURIC  ACID 


179 


insufficient,  but  it  exists  only  as  an  intermediate  product,  if  at 
all,  in  the  normal  manufacture.  Nitrous  anhydride,  also,  can 
exist  only  momentarily,  if  at  all,  as  it  decomposes  at  once  into 
nitric  oxide  and  nitrogen  dioxide  at  the  temperature  of  the 
chamber.  For  a  further  discussion  of  the  subject  see  Trautz, 
Z.  physik.  Chem.  47,  513 ;  Wentzki,  Z.  angew.  Chem.  23,  1907 ; 
Raschig,  ibid.  23.  2241,  24,  160;  Ber.  ibid.  23,  2250. 

If  it  were  possible  to  lead  into  the  chamber  pure  oxygen  and 
sulfur  dioxide,  a  small  amount  of  nitric  acid  would  convert  an 
indefinitely  large  amount  of  sulfur  dioxide  into  sulfuric  acid. 
Since,  however,  air  containing  only  21  percent  of  oxygen  mixed 
with  79  per  cent  of  nitrogen  (and  argon)  must  be  used,  there  must 


Glover 
Pyrites  Tower 

Burners 


Leaden  Chambers 
Fig.  54 


G  ay -L us  sac 
Tower 


be  a  constant  escape  of  nitrogen,  carrying  with  it  nitric  oxide 
or  nitrogen  dioxide  at  the  further  end  of  the  chamber  or  set  of 
chambers.  To  recover  these  the  gases  are  passed  through  a 
tower,  known  as  the  Gay-Lussac  tower,  Fig.  54,  in  which  they 
are  exposed  to  a  large  surface  of  concentrated  sulfuric  acid 
running  down  over  broken  coke  or  a  series  of  earthenware  plates. 
The  strong  acid  absorbs  the  oxides  of  nitrogen,  forming  nitrosyl 
sulfuric  acid,  SO2(OH)  (ONO).  This  nitrated  acid  is  then  forced 
by  compressed  air  to  the  top  of  another  tower,  called  the  Glover 
tower,  B.  Here  it  is  mixed  with  some  of  the  more  dilute 


180  A  TEXTBOOK  OF  CHEMISTRY 

acid  from  the  chamber  and  a  little  nitric  acid  to  replace  the  un- 
avoidable loss.  The  mixture  runs  down  over  broken  coke  and 
comes  in  contact  with  sulfur  dioxide  coming  from  the  pyrites 
burners,  C.  This  causes  the  denitrification  of  the  acid : 

2  SO2(OH)(ONO)  +  SO2  +  2  H2O  =  3  H2SO4  +  2  NO 

The  nitric  oxide  is,  of  course,  carried  back  into  the  first 
chamber.  When  these  towers  are  used,  only  from  25  to  40 
pounds  of  sodium  nitrate  are  required  for  the  manufacture  of 
a  ton  of  sulfuric  acid.  Without  them,  two  or  three  times  as 
much  is  required.  , 

The  acid  from  the  chambers  has  a  specific  gravity  of  1.53  to 
1.62,  and  contains  62-70  per  cent  of  the  pure  acid.  It  is  usually 
concentrated  to  about  79  per  cent  by  evaporation  in  lead  pans. 
At  this  point  the  acid  begins  to  attack  the  lead  more  strongly 
and  the  concentration  is  completed  to  a  specific  gravity  of 
1.83-1.84  and  93  to  95  per  cent,  in  glass,  platinum  or  iron,  the 
last  metal  being  only  slightly  attacked  by  the  concentrated  acid, 
although  it  dissolves  easily  in  the  dilute  acid.  If  the  concen- 
trated acid  is  distilled,  an  acid  of  constant  composition  contain- 
ing about  98.5  per  cent  of  the  pure  acid  finally  passes  over  at 
338°.  The  density  of  the  vapor  proves  that  the  process  is  not 
ordinary  boiling,  but  consists  in  the  dissociation  of  sulfuric  acid 
to  sulfur  trioxide  and  water  and  that  the  two  recombine  on 
cooling:  H2S04:£S03  +  H20 

The  specific  gravity  of  pure,  100  per  cent  sulfuric  acid  is 
slightly  less  than  that  of  a  96  to  99  per  cent  acid,  the  difference 
being  so  small  that  the  concentration  of  the  acid  cannot  be  de- 
termined satisfactorily  by  means  of  the  density. 

When  sulfuric  acid  is  mixed  with  water,  considerable  heat  is 
evolved,  and  the  volume  of  the  diluted  acid  is  considerably  less 
than  the  sum  of  the  volumes  of  the  acid  and  water  which  are 
mixed.  There  is  a  chemical  combination  between  the  acid 
and  water,  giving  a  compound  which  probably  contains  four  or 
six  hydroxyl  (OH)  groups  : 


ELECTRON  THEORY  181 

/OH 

,°H  /OH 

H 


H  VOH 

XOH 


It  is  noticeable  that  while  sulfurous  acid,          S\        ,  loses 

Ov       OH          of    XOH 
water- easily,  sulfuric  acid,      ^Sv        ,  dissociates  at  a  much 

O^    XOH 

higher  temperature  and  also  has  a  strong  tendency  to  take  up 
more  water.  Along  with  this  strong  attraction  of  the  sulfur 
atom  for  hydroxyl  groups,  which  seems  to  be  so  closely  con- 
nected with  the  addition  of  another  oxygen  atom,  is  the  fact 
that  sulfuric  acid  is  a  much  stronger  acid  than  sulfurous  acid. 
Thus  in  a  "  tenth  normal "  l  solution  of  sulfuric  acid  about  60 
per  cent  of  the  hydrogen  is  ionized,  while  in  a  tenth  normal  solu- 
tion of  sulphurous  acid,  H2SO3,  only  about  20  per  cent  is  ionized. 
The  Electron  Theory.  The  facts  which  have  just  been  given 
may  be  explained,  in  part,  by  the  electron  theory,  which  has  been 
developed  rapidly  during  the  last  few  years.  The  electron  2 
may  be  defined  as  an  atom  of  negative  electricity.  When  by 
itself  and  in  rapid  motion  its  mass  is  approximately  one  seventeen- 
hundredth  of  the  mass  of  a  hydrogen  atom.  It  is  supposed  that 
atoms  of  the  elements  are  composed,  in  part,  of  electrons  and  that 
they  may  either  gain  or  lose  these.  If  an  atom  gains  an  electron, 
it  becomes  negatively  charged ;  while  if  it  loses  one,  it  becomes 
positively  charged.  In  hydrogen  sulfide,  H2S,  it  is  supposed  that 
each  hydrogen  atom  has  lost  an  electron  which  has  been  trans- 


1  A  solution  containing  one  tenth  of  a  gram  atom  of  replaceable 
hydrogen  or  one  twentieth  of  a  gram  molecule  of  sulfuric  acid  in  one 
liter. 

2  Professor  J.  J.  Thompson  uses  the  name  "corpuscle"  instead 
of  electron.     The  evidence  of  the  existence  of  electrons  is  very 
positive. 


182  A   TEXTBOOK  OF  CHEMISTRY 

ferred  to  the  sulfur.  The  positive  hydrogen  atoms  are  then  held 
by  the  negative  sulfur  atom.  In  sulfur  dioxide  and  sulfur  triox- 
ide,  however,  the  sulfur  atom  is  supposed  to  lose  two  electrons 
to  each  oxygen  atom,  and  acquires  either  four  or  six  positive 
charges.  When  water  is  brought  in  contact  with  sulfur  dioxide, 


O  =  S  =  O,  or  sulfur  trioxide,  O  =  S^     ,  it  separates  into  a  pos- 

\) 

itive  hydrogen  atom  and  a  negative  hydroxyl  group.     The  former 
adds  itself  to  the  negative  oxygen  or  sulfur,  while  the  latter  unites 

H 


with  the  positive  sulfur.     This  gives  x+S+^T        and 

H+-~cr      +Nr 

For  the  reverse  reaction  to  occur  both 

hydrogen  and  hydroxyl  must  separate  from  the  compound,  and 
it  seems  probable  that  the  negative  hydroxyl  groups  will  be  held 
much  more  strongly  by  the  sulfur  atom  with  six  positive  charges 
than  by  the  one  which  has  only  the  equivalent  of  four.  On  the 
other  hand,  the  positive  hydrogen  might  be  expected  to  separate 
more  easily  in  the  ionic  form  from  the  sulfuric  acid  on  account  of 
the  indirect  repulsion  of  the  strongly  positive  sulfur  atom.  The 
electron  theory  is  too  recent  for  chemists  to  form  a  very  positive 
opinion  as  to  its  value,  but  it  is,  at  least,  worthy  of  careful  consid- 
eration, and  it  will  be  referred  to  repeatedly  in  the  following  pages. 

Sulfuric  Acid  as  a  Dehydrating  Agent.  On  account  of  its 
affinity  for  water,  sulfuric  acid  is  an  excellent  drying  agent  for 
all  gases  which  do  not  react  with  it  chemically.  It  is  much  more 
efficient  than  calcium  chloride.  It  will  also  take  the  elements 
of  water  from  many  such  substances  as  wood  or  sugar,  which 
contain  oxygen  and  hydrogen. 

So  much  heat  is  liberated  when  sulfuric  acid  is  mixed  with 
water  that  the  action  may  become  explosive  unless  care  is  used. 
The  concentrated  acid  should  always  be  poured  into  water  with 
which  it  is  to  be  diluted  instead  of  pouring  water  upon  the  acid. 
Why? 


SULFATES.    NORMAL  SOLUTIONS  183 

Sulfates.  Dibasic  Acids.  Either  one  or  both  of  the  hydrogen 
atoms  of  sulfuric  acid  may  be  replaced  by  a  metal,  giving  acid 
and  normal  salts,  as  acid  sodium  sulfate,  NaHSO4,  and  normal 
sodium  sulfate,  Na2SO4.  Acids  having  this  property  are  called 
dibasic.  An  acid  like  phosphoric  acid,  H3PO4,  which  forms 
three  salts  with  sodium,  NaH2PO4,  Na2HPO4  and  Na3PO4,  is 
called  tribasic.  The  basicity  depends,  however,  not  on  the 
number  of  hydrogen  atoms  in  one  molecule  of  an  acid,  but  on 
the  number  of  replaceable  hydrogen  atoms.  Thus  acetic  acid, 
C2H4O2,  is  monobasic  because  only  one  of  /ts^i^tirogen  atoms 
can  be  replaced ;  and  phosphorous  acid,  HsPOa^pjdibasic  because 
only  two  of  the  hydrogen  atoms  can  be  replaced. 

As  with  other  strong  acids,  the  normal  sulfates  of  the  metals 
of  the  sodium  and  calcium  families  are  neutral  in  reaction, 
while  the  acid  sulfates  of  all  metals  are  strongly  acid.  The  sul- 
fates of  the  metals  of  the  calcium  family,  calcium,  strontium, 
barium  and  radium,  are  difficultly  soluble  in  water,  the  solubility 
decreasing  with  increasing  atomic  weight.  Barium  sulfate  re- 
quires about  400,000  parts  of  water  for  its  solution,  while  radium 
sulfate  is  still  more  insoluble.  Lead  sulfate,  also,  is  almost 
insoluble,  but  all  other  sulfates  which  are  not  decomposed  by 
water  are  soluble.  In  general,  the  salts  of  strong  acids  are  solu- 
ble in  water,  and  this  fact  is  probably  connected  with  the  high 
degree  of  ionization  of  both  acids  and  salts.  No  explanation 
has  been  offered  for  the  exceptions  to  this  general  rule.  The 
rule  is  useful  because  we  have  to  learn  only  a  short  list  of  insolu- 
ble salts  for  these  strong  acids  and  can  then  assume  that  all  other 
salts  are  soluble. 

Normal,  Standard  and  Formular  l  Solutions.  We  have  fre- 
quently found  it  convenient  to  use  the  gram  molecule  of  sub- 
stances as  a  unit  in  dealing  with  them.  This  unit  is  often  called, 
for  the  sake  of  brevity,  one  mol.  In  working  with  acids  and 

1  The  designation  "molar"  (or  molal)  is  often  used,  but  "for- 
mular,"  if  followed  by  the  formula  of  the  substance,  is  more  definite. 
Thus  a  formular  solution  of  ferric  chloride,  FeCl3,  is  definite,  while 
a  molar  solution  of  ferric  chloride  might  refer  to  either  FeCl3  or 
Fe2Cl6. 


184  A  TEXTBOOK  OF  CHEMISTRY 

bases  it  is  often  convenient  to  use  a  gram  equivalent  instead  of  a 
gram  molecule,  as  the  unit.  The  gram  equivalent  of  an  acid 
or  base  is  that  quantity  which  is  equivalent  to  or  will  neutralize 
one  gram  molecule  of  a  monobasic  acid  or  of  a  monacid  base. 
A  solution  containing  one  gram  equivalent  of  an  acid  or  base  in 
one  liter  (or  one  milligram  equivalent  in  one  cubic  centimeter) 
is  said  to  be  normal.  Thus  a  normal  solution  of  hydrochloric 
acid  would  contain  36.47  grams  of  the  acid,  HC1,  in  one  liter ; 
but  a  normal  solution  of  sulfuric  acid  would  contain,  not  a  gram 
molecule  (98.08  grams),  but  a  gram  equivalent  (49.04  grams) 
of  sulfuric  acid,  H2SO4,  in  one  liter.  A  normal  solution  of  sodium 
hydroxide,  NaOH,  would  contain  40.01  grams  in  one  liter ;  but 
a  normal  solution  of  calcium  hydroxide,  Ca(OH)2,  if  it  could  be 
prepared,  would  contain  only  37.04  grams  per  liter.  The  advan- 
tage of  the  system  is  that  one  cubic  centimeter  of  any  normal 
solution  will  exactly  neutralize  or  be  exactly  equivalent  to  one 
milligram  equivalent  of  any  acid  or  base.  The  name  "  normal  " 
is  also  frequently  applied  to  solutions  of  salts  and  ofcother  sub- 
stances, but  such  a  use  is  liable  to  lead  to  conf usi^B  and  it  is 
better  to  call  such  solutions  "  standard  "  or^pKnular,"  a 
standard  solution  being  simply  one  whose-^roncentration  is 
known  and  a  formular  solution  one  which  contains  one  formular 
weight  in  one  liter.  The  formula  on  which  a  formular  solution 
is  based  should  always  be  given. 

If  the  term  "  normal "  is  applied  to  solutions  of  other  sub- 
stances than  acids  and  bases,  one  liter  of  the  solution  should 
always  contain  an  amount  of  the  substance  which  is  equivalent 
to  one  gram  atom  of  hydrogen  in  the  reaction  for  which  it  is 
used.  A  normal  solution  of  potassium  chloride,  KC1,  or  of 
silver  nitrate,  AgNOs,  will  contain  one  gram  molecule  of  these 
compounds  in  a  liter,  but  a  solution  of  calcium  chloride,  CaCl2, 
will  contain  only  one  half  of  a  gram  molecule.  A  solution  of 
potassium  permanganate,  KMnC>4,  will  contain  only  one  fifth  of 
a  gram  molecule,  if  to  be  used  in  an  acid  solution,  and  one  third 
of  a  gram  molecule,  if  to  be  used  in  an  alkaline  solution,  because 
one  gram  molecule  of  the  compound  will  oxidize  five  gram  atoms 


ACIDIMETRY 


185 


of  hydrogen  in  the  first  case  and  only  three  in  the  second.  A 
normal  solution  of  ferrous  sulfate,  FeSO4,  will  contain  one  gram 
molecule  of  the  compound  in  a  liter,  because  it  requires  only 
one  half  of  a  gram  atom  of  oxygen  to  oxidize  it. 

Acidimetry  and  Alkalimetry.  If  a  very  small  quantity  of  an 
indicator  (p.  122)  is  added  to  a  solution  of  hydrochloric  acid  or 
any  other  strong  acid,  on  adding  a 
solution  of  sodium  hydroxide  or  some 
other  strong  base,  the  change  in 
color  of  the  indicator  will  show  very 
sharply  when  the  acid  has  been  ex- 
actly neutralized  by  the  base. 

If  we  have  a  normal  solution  of 
hydrochloric  acid  (containing  36.47 
milligrams  of  the  acid,  HC1,  in  1  cc., 
as  defined  above),  it  is  easy  by 
measuring  this  from  a  burette  (Fig. 
55)  exactly  to  neutralize  a  solution 
containing^  strong  base.  The  num- 
ber of  cubi(£centimenters  of  the  acid 
used  will  give  at  once  the  numbers 
of  milligram  equivalents  of  the  base 
which  were  present  in  the  solution 
neutralized.  Thus  one  cubic  centi- 


r 

D 

:  ».                         f 

j 

™ 

< 

1 

I* 

• 

r 

Fig.  55 


meter  of  normal  hydrochloric  acid  will  exactly  neutralize  40.01 
milligrams  of  sodium  hydroxide,  NaOH,  56.11  milligrams  of 
potassium  hydroxide,  KOH,  or  37.04  milligrams  of  calcium 
hydroxide,  Ca(OH)2.  In  the  same  manner,  by  means  of  a 
normal  solution  of  potassium  hydroxide,  KOH  (containing 
56.11  milligrams  in  1  cc.),  the  number  of  milligram  equivalents 
of  any  strong  acid  contained  in  a  given  solution  can  be  readily 
determined.  The  process  of  making  such  determinations  is 
called  acidimetry  or  alkalimetry.  The  choice  of  an  indicator 
and  th*e  application  of  the  process  to  some  cases  involving 
weak  acids  and  bases  will  be  discussed  in  a  later  chapter 
(p.  387). 


186  A  TEXTBOOK  OF  CHEMISTRY 

Pyrosulfates.  When  acid  sodium  sulfate  is  heated,  it  loses 
water  and  is  converted  into  a  salt  which  is  called,  for  this  reason, 
sodium  pyrosulfate : 

2  NaHSO4  =  Na2S207  +  H2O 

Sodium 
Pyrosulfate 

A  solution  of  sulfur  trioxide  in  sulfuric  acid  doubtless  always 
contains  pyrosulfuric  acid,  H2S2O7,  and  pure  pyrosulfuric  acid 
is  a  definite  compound  which  melts  at  35°,  but  it  is  very  unstable, 
dissociating  easily  into  sulfur  trioxide  and  sulfuric  acid.  In 
solution  the  pyrosulfates  take  up  water  and  pass  back  into  acid 
sulfates. 

Hyposulfites.  When  zinc  is  dissolved  in  a  solution  of  sulfurous 
acid,  H2SO3,  the  acid  is  reduced  and  zinc  hyposulfite  is  formed  : 

Zn  +  2  H2SO3  =  ZnS2O4  +  2  H2O 

Zinc 
Hyposulfite 

Zinc  hyposulfite  is  a  salt  of  an  unstable  acid,  hyposulfurous 
acid,  H2S2O4,  which  is  not  known  in  the  free  state.  The  salts 
are  very  quickly  oxidized  to  sulfites  in  the  air  and  are  powerful 
reducing  agents.  Sodium  hyposulfite,  Na2S2O4,  is  manufactured 
for  use  in  the  reduction  of  indigo  to  indigo  white  (p.  341). 

Hyposulfurous  acid  and  the  hyposulfites  must  not  be  con- 
fused with  thiosulfuric  acid  and  thiosulfates  which  were  formerly 
called  by  the  same  name  (see  the  next  paragraph).  Some 
authors  prefer  to  call  the  acid  hydrosulfurous  acid  to  avoid 
possible  confusion.  Neither  acid  corresponds  to  the  formula 
(H2SO2),  which  we  should  logically  expect  for  a  hyposulfurous 
acid. 

Thiosulfates.  A  solution  of  sodium  sulfite,  Na2SO3,  will 
dissolve  sulfur,  and  there  may  then  be  crystallized  from  the  solu- 
tion a  salt  called  sodium  thiosulfate,  Na2S2O3.5  H2O.  The 
change  is  similar  to  the  oxidation  of  sodium  sulfite  to*  sodium 
sulfate :  Na2SO3  +  O  =  Na2SO4 

Na2SO3  +  S  =  Na2S2O3 


PERSULFURIC  ACID  187 

one  atom  of  sulfur  taking  the  place  of  an  atom  of  oxygen,  and  the 
name  thiosulfatc  (from  Greek  Odov,  sulfur)  is  given  to  the 
salt  for  this  reason.  The  salt  has  been  long  known  and  was 
originally  called  sodium  hyposulfite,  a  name  which  still  clings  to 
it  among  druggists  and  photographers.  It  is  extensively  used 
in  photography  as  a  solvent  for  silver  chloride  or  bromide  in 
"  fixing  "  pictures. 

If  a  solution  of  a  thiosulfate  is  acidified,  the  thiosulfuric  acid 
at  first  liberated  decomposes  with  the  liberation  of  sulfur  : 

H2S203  =  S02  +  H20  +  S 

Iodine  converts  sodium  thiosulfate  into  sodium  tetrathionate  : 
2  Na2S2O3  +  I2  =  Na2S406  +  2  Nal 

Sodium 
Tetrathionate 

This  reaction  is  much  used  in  connection  with  standard  iodine 
solutions,  in  volumetric  analysis. 

Persulfuric  Acid.  When  a  solution  of  acid  potassium  sulfate, 
KHSO4,  is  electrolyzed  with  a  high  current  density,  that  is,  with 
a  current  strong  in  comparison  with  the  surface,  at  the  anode, 
as  the  anions,  HSO4~,  are  discharged  they  combine,  in  part,  with 
other  anions  of  the  same  kind  to  form  persulfuric  acid, 
H— O— SO2— O— O— SO2— OH  or  H2S2O8.  The  persulfuric  acid 
then  reacts  with  some  of  the  acid  potassium  sulfate  present  to 
form  potassium  persulfate,  which  is  rather  difficultly  soluble: 

2  KHSO4  +  H2S2O8  =  K2S208  +.2  H2SO4 

Persulfuric  acid  is  also  formed  when  hydrogen  peroxide,  H2O2, 
is  added  to  concentrated  sulfuric  acid : 


H— O— SO2—  iO-H  H|— O— O— !H  H— Oj— SO2— OH 
or  2  H2SO4  +  H2O2  =  H2S2O8  +  2  H2O 

Persulfuric  acid  and  the  persulfates  are  used  as  oxidizing 

agents. 


188  A  TEXTBOOK  OF  CHEMISTRY 

*  Permonosulfuric  Acid.     When  a  solution  of  persulfuric  acid 
is  diluted  and  allowed  to  stand,  it  changes  to  permonosulfuric 
acid: 

H— O— SO2— O— O— SO2-OH  +  HOH 

=  H— O— S02— O— OH  +  H2S04 
or  H2S2O8  +  H2O  =  H2SO5  +  H2SO4 

The  solution  was  formerly  known  as  Caro's  acid  and  is  some- 
times used  as  an  oxidizing  agent  for  organic  compounds.  For 
instance,  it  will  oxidize  aniline  to  nitrobenzene. 

*  Polythionic  Acids.     A  series  of  acids  having  from  two  to  six 
atoms  of  sulfur  in  a  molecule  has  been  obtained.     These  are : 

Dithionic  acid          H2S2Oe 

Trithionic  acid         H^SaOe 

Tetrathionic  acid     H2S4O6.     (See  thiosulfuric  acid,  above.) 

Pentathionic  acid    H2S5O6 

Hexathionic  acid      H2S6O6 
These  acids  need  not  be  considered  in  detail  here. 

Compounds  of  Sulfur  containing  Halogens.  Quite  a  num- 
ber of  such  compounds  are  known.  All  of  them  except  sulfur 
hexafluoride,  SF6,  are  hydrolyzed  by  water,  giving  hydrochloric 
acid  or  a  halogen  acid  and  some  acid  of  sulfur.  Thus  sulfuryl 
chloride,  SO2C12,  gives : 

HOH         /OH 

QfV/ 

HOH 


+  =  SO2<         +  2  HC1 

X)H 


*  Sulfur  Monochloride,  S2C12,  is  a  clear,  amber-colored  liquid 
formed  by  passing  chlorine  over  heated  sulfur.  It  boils  at  138° 
and  is  hydrolyzed  by  water  to  hydrochloric  acid,  thiosulfuric 
acid  and  sulfur.  Its  specific  gravity  is  1.7055. 

Because  of  the  strong  affinities  of  sulfur  for  oxygen  and  of 
chlorine  for  metals,  the  chlorides  of  a  number  of  metals,  which 
it  is  difficult  to  prepare  from  the  oxides  otherwise,  may  be  ob- 
tained by  passing  a  mixture  of  sulfur  monochloride  and  chlorine 
over  the  heated  oxides.  (E.  F.  Smith.)  It  is  also  used  in  the 
manufacture  of  India  rubber, 


SELENIUM  189 

*  Chlorosulfonic  Acid 1  is  easily  prepared  by  passing  hydro- 
chloric acid  gas  through  warm,  fuming  sulfuric  acid. 

Ov  O^    XC1 

>S=0  +  HC1  =     lS< 
CT  CT      X)H 

It  boils  at  152-153°  and  is  easily  hydrolyzed  by  water.  Its 
specific  gravity  is  1.766  at  18°. 

*  Sulfuryl  Chloride,  SO2C12,  may  be  prepared  by  the  union 
of  sulfur  dioxide  and  chlorine,  or,  more  easily,  by  boiling  chloro- 
sulfonic  acid  with  mercuric  sulfate,  which  acts  as  a  catalytic 
agent,  causing  it  to  decompose  in  accordance  with  the  equation  : 

2  SO2OHC1  =  H2SO4  +  SO2C12 

Sulfuryl  chloride  is  sometimes  called,  less  correctly,  the  chlo- 
ride of  sulfuric  acid.  Similar  acid  chlorides  may  be  formed  by 
replacing  the  hydroxyl  of  other  acids  with  chlorine.  Acid  chlo- 
rides are  hydrolyzed  by  water,  giving  the  acid  from  which  they 
are  derived  and  hydrochloric  acid. 

Sulfuryl  chloride  boils  at  69.1°  and  has  a  specific  gravity  of 

1. 6674  at  ^-2. 

Selenium,  Se,  79.2,  is  found  in  small  amounts  as  selenides  of 
metals,  associated,  usually,  with  sulfides  of  these  same  metals. 
When  such  sulfides  are  used  for  the  manufacture  of  sulfuric 
acid,  selenium  is  sometimes  found  in  the  dust  flues  of  the  pyrites 
burners  and  in  the  slime  on  the  bottom  of  lead  chambers.  Both 
selenium  and  tellurium  are  found  in  considerable  quantities  in 
the  slimes  from  electrolytic  copper  refining. 

It  occurs  in  several  allotropic  forms,  the  red  variety  obtained 
by  crystallization  from  carbon  bisulfide  and  a  gray  metallic 
form  obtained  by  melting  either  of  the  other  forms  being  the 
best  defined.  An  amorphous  form  is  also  known.  The  metallic 
form  melts  at  217°  and  has  a  specific  gravity  of  4.8.  Selenium 
boils  at  680°. 

1  Acids  containing  the  group  S02OH  are  called  sulfonic  acids. 

2  This  means  the  specific  gravity  at  20°  referred  to  water  at  4°. 


190  A  TEXTBOOK  OF  CHEMISTRY 

The  metallic  form  of  selenium  conducts  electricity.  Its  con- 
ductivity is  very  greatly  affected  by  changes  of  temperature  or 
by  exposure  to  light,  and  several  important  applications  of  this 
property  have  been  invented,  one  of  the  most  important  being 
in  stellar  photometry. 

Hydrogen  Selenide,  H2Se,  may  be  prepared  by  the  action  of 
hydrochloric  acid  on  ferrous  selenide,  FeSe.  It  is  very  poison- 
ous, and  the  odor  is  more  unpleasant  than  that  of  hydrogen  sul- 
fide.  Compare  the  series,  water,  hydrogen  sulfide,  hydrogen 
selenide,  in  this  respect.  Berzelius,  one  of  the  early  workers 
with  hydrogen  selenide,  reports  that  after  breathing  a  single 
bubble  of  the  gas  he  so  far  lost  the  sense  of  smell  for  several 
hours  that  he  could  not  distinguish  the  odor  of  strong  ammonia. 

Selenium  Dioxide,  SeO2,  is  a  white  solid  prepared  by  burning 
selenium  in  a  current  of  oxygen.  It  gives  selenious  acid,  H2SeO3, 
on  solution  in  hot  water.  From  a  solution  of  selenious  acid 
sulfur  dioxide  precipitates  selenium  as  a  red  powder : 

H2SeO3  +  2  SO2  +  H2O  =  Se  +  2  H2SO4 

Selenic  acid,  H2SeO4,  is  formed  by  the  action  of  bromine  on 
silver  selenite,  the  silver  bromide  formed  separating  as  a  precipi- 
tate * 

Ag2SeO3  +  Br2  +  H2O  =  H2SeO4  +  2  AgBr 

Selenic  acid  loses  oxygen  easily  and  is  a  strong  oxidizing  agent. 

Tellurium,  Te,  127.5,  is  found  in  combination  with  gold, 
silver,  copper  and  bismuth.  It  is  a  white,  metallic-looking 
solid,  which  melts  at  452°,  boils  at  1400°  and  has  a  specific 
gravity  of  6.44.  Its  most  interesting  compounds  are  hydrogen 
telluride,  H2Te,  tellurium  dioxide,  TeO2,  tellurium  trioxide, 
TeO3,  tellurous  acid,  H2TeO3,  and  telluric  acid,  H2TeO4. 

Atomic  Weight  of  Tellurium.  Very  many  determinations  of 
the  atomic  weight  of  tellurium  have  given  values  about  127.5, 
decidedly  higher  than  the  atomic  weight  of  iodine,  126.9.  The 
properties  of  tellurium  and  especially  the  formulas  of  its  com- 
pounds indicate  that  it  should  precede  iodine  in  the  Periodic 
System,  and  this  has  led  to  many  attempts  to  determine  whether 


TELLURIUM.    GROUP  VI  191 


the  material  used  for  the  atomic  weight  determinations  has  been 
pure,  or  whether,  possibly,  it  may  have  contained  some  other 
element  from  which  it  is  unusually  difficult  to  separate  a  pure 
tellurium  compound.  Some  of  these  attempts  to  discover  a 
method  of  preparing  tellurium  of  greater  purity  and  lower  atomic 
weight  have  seemed,  for  a  time,  to  be  successful ;  but  none  of 
these  lower  results  for  the  atomic  weight  has  been  confirmed  by 
other  workers,  and  it  seems  pretty  certain  that  the  atomic  weight 
of  tellurium  is  greater  than  that  of  iodine.  See  Browning  and 
Flint,  Z.  anorg.  Chem.  64,  104,  112,  and  68,  251 ;  and  Harcourt 
and  Baker,  J.  Chem.  Soc.  100,  1311. 

General  Properties  of  the  Elements  of  the  Sixth  Group.  Just 
as  chlorine,  bromine  and  iodine  are  much  more  closely  related 
in  their  properties  than  fluorine  is  related  to  them,  oxygen  stands 
somewhat  by  itself  in  the  sixth  group,  while  the  relationships 
between  sulfur,  selenium  and  tellurium  are  comparatively  close. 
The  halogens  have  a  valence  of  one  in  their  compounds  with 
hydrogen,  as  in  HF,  HC1,  etc.,  and  a  maximum  valence  of  seven  in 
their  compounds  with  oxygen,  as  in  C^O?,  HCIO^H— O— ClOa), 
HIO4,  etc.  The  elements  of  the  sulfur  family  have  a  valence  of 
two  in  their  compounds  with  hydrogen,  as  in  H2O,  H2S,  etc.,  and 
a  maximum  valence  of  six  toward  oxygen,  as  in  SOs,  £[2804, 

H-CX       ,0 

>SC    ,H2TeO4,  etc. 
H-  (X     ^O 

There  is  a  similar  gradation  of  physical  properties  in  the  two 
groups :  fluorine  and  chlorine  are  gases,  bromine  a  liquid  and 
iodine  a  solid,  with  increasing  depth  of  color  as  the  atomic  weight 
increases.  In  the  same  way,  oxygen  is  colorless  (ozone  is  blue), 
sulfur  is  a  light  yellow  solid,  selenium  is  dark  red,  and  tellurium 
is  opaque  and  has  many  of  the  properties  of  a  metal.  Indeed, 
if  it  were  not  for  its  position  in  the  Periodic  System  and  the  re- 
semblance between  the  formulas  of  its  compounds  and  those  of 
selenium  and  sulfur,  tellurium  would  be  classed  as  a  metal,  or, 
at  least,  as  a  half  metal.  But  it  will  be  seen  that  in  the  succes- 
sive groups  the  metallic  properties  become  more  and  more 


192  A  TEXTBOOK  OF  CHEMISTRY 

marked  with  increasing  atomic  weight.  Thus  arsenic,  antimony 
and  bismuth,  of  the  fifth  group,  are  usually  classed  as  metals, 
though  all  of  them  are  brittle.  Tin  and  lead,  of  the  fourth  group, 
are  clearly  metals  and  are  malleable,  though  deficient  in  tenacity. 
The  most  typical  compounds  of  the  sixth  group  are  the  follow- 

mg:         H2O  H2S  H2Se  H2Te 

O3  SO2  SeO2  TeO2 

SO3  TeO3 

H2S03  H2Se03  H2Te03 

H2SO4  H2SeO4  H2TeO4 

As  in  the  halogen  family,  the  chemical  activity,  in  general, 
decreases  with  increasing  atomic  weight.  Hydrogen  sulfide 
dissociates  at  a  much  lower  temperature  than  water  does,  and 
sulfur  dioxide  will  take  oxygen  from  selenious  acid,  reducing 
it  to  free  selenium. 

As  manganese  forms  compounds  which  resemble  some  of  the 
compounds  of  chlorine,  there  are  four  metals  of  the  sixth  group, 
chromium,  molybdenum,  tungsten  (symbol  W,  from  wolfram) 
and  uranium,  which  form  oxides  and  salts  of  acids  similar  to  the 
oxides  and  acids  of  sulfur.  The  oxides  are :  CrO3,  MoO3,  WO3, 
UO3 ;  and  the  corresponding  sodium  salts  of  the  acids  are : 
Na2CrO4,  Na2MoO4,  Na2WO4.  Uranium  forms  a  compound, 
UO2(OH)2,  similar  in  composition  to  sulfuric  acid,  but  it  is  a 
base  rather  than  an  acid,  another  illustration  of  the  fact  that  an 
increase  in  the  atomic  weight,  within  a  given  group,  increases 
the  metallic  properties  of  the  element. 

Crystals.  When  substances  solidify  from  the  molten  condi- 
tion or  when  they  separate  on  the  evaporation  of  a  solution, 
molecules  of  the  same  kind  frequently  arrange  themselves  in 
definite,  geometrical  relations  to  each  other,  forming  solids 
bounded  by  plane  faces,  which  are  called  crystals.  This  prop- 
erty has  already  been  mentioned  as  an  important  means  for 
preparing  pure  substances.  It  is  also  a  very  important  and  char- 
acteristic property  of  individual  substances,  and  the  shapes  of 
the  crystals  of  different  compounds  offer  such  an  infinite  variety 


CRYSTALS 


193 


that  they  may  frequently  be  used  as  a  very  positive  means  of 
identification. 

In  spite  of  the  large  number  of  crystalline  forms  all  crystals 
may  be  classified  in  six  systems.  These  systems  are  most  easily 
defined  by  referring  each  to  axes,  which  are  used  in  much  the 
same  manner  as  the  coordinates  of  analytical  geometry,  to  de- 
fine the  structure  of  the  crystal  and  the  relation  between  the 
planes  bounding  its  surface. 

These  systems  are : 

1.  The  Isometric  or  Regular  System.  Three  equal  axes  at 
right  angles.  Some  of  the  simplest  forms  of  this  system  are  the 


Fig.  56  Fig.  57  Fig.  58 

cube  (Fig.  56),  octahedron  (Fig.  57),  rhombic  dodecahedron  (Fig. 
58)  and  the  tetrahedron  (Fig.  59).    The  last  has  only  half  of  the 


Fig.  59  Fig.  60 

faces  of  the  octahedron  and  is  called  a  hemihedral  form.  Com- 
binations of  two  or  more  forms  are  also  common.  Figure  60  is 
a  combination  form  called  a  tetrahexahedron. 


194 


A  TEXTBOOK  OF  CHEMISTRY 


2.  The  Tetragonal  System.  Three  axes  at  right  angles,  two 
of  them,  only,  being  equal.  The  tetragonal  pyramid  (Fig.  61) 
and  the  square  prism  (Fig.  62)  are  common  forms.  Figure  63 
shows  a  combination  of  the  two. 


Fig.  61 


Fig.  62 


3.  The  Rhombic  System.  Three  axes  at  right  angles  but  of 
unequal  length.  Rectangular  and  rhombic  prisms  (Figs.  64  and 
65)  and  pyramids  are  illustrations. 


'1 

^ 

~1 

1 

f^L 

1 

•         m 

'r. 


Fig.  64 


Fig.  65 


4.  The  Hexagonal  System.     Three  axes  in  the  same  plane  at 
an  angle  of  60°  with  each  other  and  a  fourth  at  right  angles  to 


CRYSTALS 


195 


the  plane  of  the  other  three.  Common  forms  are  the  hexagonal 
pyramid  (Fig.  66)  and  prism  (Fig.  67)  and  the  rhombic  hexahe- 
dron (Fig.  68),  which  has  only  half  of  the  faces  of  the  pyramid. 


Fig.  66 


Fig.  67 


5.  The  Monoclinic  System.  Two  axes  at  right  angles  and  a 
third  at  right  angles  to  one  and  inclined  to  the  other,  the  three 
axes  being  unequal  (Fig.  69). 


Fig.  68 


196 


A  TEXTBOOK  OF  CHEMISTRY 


Fig.  70 


6.  The  Triclinic  System.  Three  unequal  axes,  all  inclined 
(Fig.  70). 

Crystals  rarely  exhibit  the  complete  geometrical  forms  which 
are  the  ideal  to  which  they  are  referred.  They  always  have  an 

internal  structure  characteris- 
tic of  these  forms,  however, 
and  this  can  often  be  detected 
by  their  optical  properties. 
Thus  for  the  regular  system 
light  travels  with  the  same 
velocity  in  all  directions 
through  the  crystal.  For 
other,  less  symmetrical  forms, 
the  velocity  is  different  in 
different  directions  and  this 
causes  double  refraction,  polarization  of  light  and  other  phe- 
nomena frequently  used  for  the  identification  of  the  crystal  form. 
The  angles  between  the  faces  of  crystals  are  also  accurately 
fixed  by  the  system  to  which  they  belong  and  the  properties 
of  the  individual  substance,  and  the  measurement  of  these  angles 
is  used  for  purposes  of  identification. 

Such  substances  as  sulfur,  which  crystallize  in  two  different 
forms,  are  called  dimorphous. 

Different  substances  which  crystallize  in  the  same  form  are 
called  isomorphous.  A  crystal  of  a  substance  should  grow  if 
placed  in  a  supersaturated  solution  of  a  substance  with  which 
it  is  isomorphous. 

EXERCISES 

1.  One  liter  of  water  at  0°  absorbs  4.37  volumes  of  hydrogen  sul- 
fide.     What  part  of  a  gram  molecule  of  the  gas  does  the  solution  con- 
tain ?     What  would  be  the  theoretical  depression  of  the  freezing  point 
if  the  compound  were  completely  ionized  to  HS~  and  H+  ?    If  it  were 
not  at  all  ionized  ? 

2.  How  many  liters  of  air  containing  21  per  cent  of  oxygen  will  be 
required  to  burn  one  liter  of  hydrogen  sulfide  to  water  and  sulfur  ?   How 
many  liters  would  be  required  to  burn  it  to  sulfur  dioxide  and  water  ? 


SULFUR,  SELENIUM  AND  TELLURIUM  197 


3.  How  many  cubic  feet  of  air  at  0°  (1  cu.  ft.  =  28.315  liters)  will  be 
required  to  burn  enough  pyrites  to  make  one  ton  of  chamber  acid  of 
70  per  cent?     (1  ton  =  907.18  kilograms.)     How  many  cubic  feet  of 
air  must  be  introduced  into  the  chamber  to  convert  the  sulfur  dioxide 
to  sulf uric  acid  ? 

4.  Solve  the  same  problem,  substituting  the  metric  ton  and  cubic 
meters  for  ton  and  cubic  feet.     How  many  kilograms  of  water  must  be 
introduced  in  the  chamber  ? 

5.  What  is  the  percentage  increase  in  the  volume  of  the  air  which 
must  be  used  if  the  temperature  is  25°  instead  of  0°  ? 

6.  How  many  grams  of  ferric  oxide  will  be  obtained  by  burning  one 
kilogram  of  iron  pyrites  ? 


CHAPTER  XII 
NITROGEN 

SYMBOL,  N.    ATOMIC  WEIGHT,  14.01. 

Occurrence  and  Natural  History  of  Nitrogen.  Fluorine  and 
oxygen,  the  first  elements  of  the  seventh  and  sixth  groups,  are 
very  active ;  and  while  oxygen  is  found  free  in  the  air,  this  seems 
to  be  more  on  account  of  its  abundance  and  because  nearly  all  of 
the  other  elements  in  nature  are  already  combined  with  oxygen 
than  because  of  any  lack  of  activity.  Nitrogen,  the  first  element 
of  the  fifth  group,  in  very  striking  contrast  to  oxygen  and  fluorine, 
is  found  chiefly  in  the  free  state  in  the  atmosphere  because  it 
does  not  readily  combine  with  any  of  the  other  elements  avail- 
able in  the  earth. 

Until  1894  it  was  supposed  that  the  gas  remaining  when  oxy- 
gen, carbon  dioxide  and  moisture  were  removed  from  air  was 
pure  nitrogen.  Rayleigh  and  Ramsay  showed  at  that  time  that 
the  residue  left  after  the  removal  of  these  substances  still  con- 
tained about  1.2  per  cent  of  gases,  chiefly  argon  (p.  235),  which 
are  even  more  inert  than  nitrogen.  Nitrogen  forms,  however, 
nearly  78  per  cent  of  the  volume  of  dry  air ;  and  as  the  air  above 
a  square  centimeter  of  the  earth's  surface  weighs  about  one 
kilogram  or  that  above  a  square  meter  weighs  more  than  ten 
tons,  it  is  evident  that  the  amount  of  nitrogen  in  the  atmosphere 
is  very  large,  though  it  is  small  in  comparison  with  the  amounts 
of  those  elements  which  make  up  the  bulk  of  the  solid  crust 
of  the  earth.  In  combination,  nitrogen  is  found  in  all  living 
organisms  as  an  essential  constituent,  the  four  most  important 
elements  in  organic  matter  being  carbon,  nitrogen,  hydrogen  and 
oxygen.  But  while  plants  can  obtain  the  carbon  for  their  growth 
from  the  carbon  dioxide  of  the  air  and  the  oxygen  and  hydrogen 
from  the  moisture  of  the  soil,  very  few,  if  any,  of  the  higher 
forms  of  plant  life  can  use  the  nitrogen  of  the  air  directly.  The 

198 


NITROGEN 


199 


larger  part  of  the  nitrogen  which  is  essential  for  the  growth  of 
crops  must  be  supplied  in  the  form  of  compounds  which  result 
from  the  decay  of  animal  or  vegetable  substances,  or  from  the 
combination  of  oxygen  and  nitrogen  in  the  air  through  electrical 
agencies.  There  are,  however,  a  few  plants,  especially  clover, 
alfalfa  and  other  leguminous  plants,  which  are  able  to  assimilate 
the  nitrogen  t>f  the  air  with  the  aid  of  bacteria  which  grow  in 
nodules  on  their  roots. 

The  decay  of  organic  matter  containing  nitrogen  is  always 
caused  by  the  growth  of  bacteria.  In  the  absence  of  air,  the 
conditions  of  decomposition  lead  to  the  reduction  of  the  nitrogen 
to  ammonia.  In  the  presence  of  air,  as  in  a  well-aerated  soil, 
the  nitrifying  bacteria,  which  are  usually  present,  will  convert 
the  nitrogen  to  nitric  acid,  HNOs,  which  generally  finds  enough 
potassium,  calcium  or  sodium  present  to  form  saltpeter,  KNOs, 
calcium  nitrate,  Ca(NO3)2  or  sodium  nitrate,  NaNOa.  All 
plants  can  readily  assimilate  the  nitrogen  of  the  nitrates  and  so 
it  finds  its  way  back  into  the  organic  compounds  of  the  plant 
life.  As  nitrogen  in  a  readily  available  form  is  essential  to  the 
growth  of  wheat,  corn  and  other  crops,  sodium  nitrate,  the 
cheapest  of  the  commercial  nitrates,  and  ammonium  sulfate, 
(NH4)2SO4,  also  a  comparatively  cheap  nitrogen  compound,  are 
often  used  as  fertilizing  materials. 

The  course  of  nitrogen  in  nature  is  sh<5wn  by  the  following 
diagram : l 

Leguminous  plants  with 


the  help  of  bacteria 
Atmospheric 

elftntriftitv                          Plants                      ' 

Atmosphe 
Nitrogei 

_:„                          "%^ 

Nitrates 

fing 

i 

Nitrif 

Organic  compounds 
of  Nitrogen 

Denitrifying 
bacteria 

Ni' 
bac 

ktrif: 
teris 

Decay 
and  ani 
or  dis, 

of  plant 
mal  tissues 
tillation 

Mitr 

ying  [—  • 
Ammonia 

See  Abegg,  Handbuchderanorg.  Chemie,  Bd.  3,  Abth.  3,  S.  215. 


200  A  TEXTBOOK  OF  CHEMISTRY 

Preparation  and  Properties  of  Nitrogen.  Nitrogen  which  is 
pure  with  the  exception  of  about  1.2  per  cent  of  the  inert  gases 
of  the  argon  family  may  be  prepared  by  burning  phosphorus 
in  air  or  by  passing  air  over  heated  copper  turnings,  which 
will  take  up  the  oxygen.  A  very  convenient  method  for  the 
preparation  of  considerable  quantities  of  nitrogen  is  to  pass  a 
mixture  of  air  and  hydrogen  through  a  tube  containing  hot 
copper  oxide.  By  keeping  the  hydrogen  slightly  in  excess  of 
the  amount  necessary  to  combine  with  the  oxygen,  a  part  of  the 
copper  oxide  will  be  reduced  to  metallic  copper ;  and  it  is  then 
easy  to  regulate  the  currents  of  the  two  gases  so  that  both 
copper  and  copper  oxide  will  be  present  in  the  tube.  The  hy- 
drogen must,  of  course,  be  mixed  with  the  air  at  the  point  where 
the  two  gases  come  in  contact  with  the  copper.  (Why  ?) 

To  prepare  nitrogen  free  from  argon  some  compound  of  nitro- 
gen must  be  used,  ammonium  nitrite,  NH4NO2,  being  most 
suitable.  When  a  solution  of  the  salt  is  warmed,  it  decomposes 
to  water  and  nitrogen : 

NH4NO2  =  N2  +  2  H2O 

Instead  of  ammonium  nitrite  a  mixture  of  sodium  nitrite  and 
ammonium  chloride  may  be  used,  the  following  reversible  reac- 
tion occurring  first : 

NaNO2  +  NH4C1  ^±  NH4NO2  +  NaCl 

Nitrogen  is  a  colorless,  odorless  gas  which  condenses  to  a 
liquid  that  boils  at  —  196°  and  freezes  at  —  210°.  It  is  very 
inert  and  does  not  combine  with  any  element  at  ordinary  tem- 
peratures, except  under  the  influence  of  microorganisms,  as 
referred  to  above.  At  the  high  temperatures  of  the  electric 
discharge  it  combines  with  oxygen  to  form  nitric  oxide,  NO, 
and  this  fact  has  been  used  recently  as  a  basis  for  the  commercial 
manufacture  of  nitrates  (see  below).  As  the  supply  of  sodium 
nitrate  in  Chile,  the  only  large  supply  now  known,  will  be  ex- 
hausted within  a  comparatively  few  years,  it  seems  certain  that 
this  manufacture  is  destined  to  be  very  important. 


AMMONIA 


201 


At  moderate  temperatures  and  under  high  pressure  nitrogen 
and  hydrogen  combine  to  form  ammonia,  NH3,  but  the  reaction 
is  very  slow  without  some  catalytic  agent. 

The  combination  is  exothermic  and  hence  the  equilibrium, 


is  shifted  toward  the  decomposition  of  the  ammonia  at  high 
temperatures.  As  the  volume  decreases  as  the  gases  combine, 
pressure  shifts  the  equilibrium  to  the  right.  Principle  of  van't 
Hoff-Le  Chatelier,  p.  111.  (See  Haber,  Z.  Elektrochem.  16, 
244.)  A  careful  study  of  the  conditions  best  suited  for  the 
reaction  has  proved  so  encouraging  that  the  Badische  Anilin 
Soda  Fabrik  in  Germany  is  preparing  for  the  manufacture  of 
synthetic  ammonia  on  a  large  scale.  The  best  catalyzers  for 
the  reaction  seem  to  be  metallic  osmium  or  uranium.  (See  J. 
Ind.  and  Eng.  Chem.  5,  328  (1913).) 

Several  metals,  especially  lithium,  magnesium  or  calcium,  com- 
bine with  nitrogen  to  form  nitrides,  at  high  temperatures : 

,Li 


6Li 


or 


Li 


2Li3N 

Lithium 
Nitride 


=  Mg3N2 
Magnesium 
Nitride 


Nitrogen  will  not  support  combustion  nor  burn. 

Ammonia.  When  organic  matter  containing' nitrogen  decom- 
poses with  exclusion  of  air,  either  under  the  influence  of  bacteria 
or  of  heat,  the  nitrogen  is  converted  partly  into  ammonia,  NH3. 
In  this  way  ammonia  is  always  found  in  sewage  or  in  piles  of 
manure.  It  is  also  formed  in  the  destructive  distillation  of  bi- 
tuminous coal  for  the  manufacture  of  illuminating  gas.  The 
aqueous  portion  of  the  liquid  distillate  from  the  coal  furnishes, 
at  present,  the  chief  source  of  the  ammonia  of  commerce.  These 


202  A  TEXTBOOK  OF  CHEMISTRY 

ammoniacal  gas  liquors  are  mixed  with  slaked  lime  and  distilled, 
the  lime  retaining  sulfur  and  other  impurities.  The  distillate 
is  mixed  with  hydrochloric  or  sulfuric  acid  and  evaporated  to 
obtain  ammonium  chloride,  NH4C1,  (NH3  +  HC1),  or  ammo- 
nium sulfate,  (NH4)2SO4,  (2  NH3  +  H2SO4). 

From  these  salts  the  ammonia  may  be  liberated  by  any  strong 
base,  as  sodium  hydroxide  or  calcium  hydroxide : 
(NH4)2SO4  +  Ca(OH)2^±CaSO4  +  2  NH4OH  ^±  2  NH3  +  2  H2O 
Ammonia  may  also  be  obtained  by  hydrolyzing  a  nitride  with 
water : 

Li3N  +  3  HOH  =  NH3  +  3  LiOH 

For  laboratory  or  lecture  purposes  ammonia  gas  is  most  easily 
obtained  by  boiling  a  strong  solution  known  as  aqua  ammonia  and 
passing  the  gas  through  a  cylinder  filled  with  quicklime  to  dry  it. 

Properties  of  Ammonia.  Ammonia  is  a  colorless  gas  with  a 
very  pungent  odor.  It  is  very  easily  soluble  in  water,  and  hence 
must  be  collected  by  displacement  of  air  (should  the  mouth  of 
the  bottle  point  up  or  down  ?)  or  over  mercury.  Water  at  0° 
absorbs  about  1000  times  its  volume  of  the  gas,  but  gives  off  a 
large  part  of  it  on  warming  gently  and  all  of  it  on  boiling.  The 
density  of  the  solution  is  less  than  that  of  water,  a  28  per  cent 
solution  having  a  specific  gravity  of  0.90. 

Ammonia  combines  directly  with  acids  to  form  ammonium 
salts,  in  which  the  hydrogen  of  the  acid  combines  with  the  am- 
monia to  form  the  ammonium  group,  NH4,  a  radical  which  in 
its  compounds  possesses  properties  very  closely  resembling  the 
properties  of  potassium  or  sodium  : 

NH3  +  HC1  =  NH4C1 

Ammonium 
Chloride 

2  NH3  +  H2S04  =  (NH4)2SO4 

Ammonium 

Sulfate  .1 

These  compounds  are  most  satisfactorily  explained  by  sup- 
posing that  nitrogen  is  trivalent  when  combined  exclusively  with 


AMMONIA  203 

hydrogen  or  with  positive  groups  but  may  become  quinquivalent 
when  one  of  the  groups  or  atoms  is  negative  : 


Any  acid  may  take  up  as  many  molecules  of  ammonia  as  it 
has  of  replaceable  hydrogen  atoms.  Thus  a  monobasic  acid, 
as  nitric  acid,  HNOa,  may  combine  with  one  molecule  of  am- 
monia, forming  ammonium  nitrate,  NH^NOs,  or  a  tribasic 
acid,  as  H3PO4,  may  combine  with  three  molecules. 

The  formation  of  an  ammonium  salt  may  be  very  prettily 
illustrated  by  filling  two  cylinders  with  ammonia  and  hydro- 
chloric acid  gas  respectively.  On  bringing  the  mouths  of  the 
cylinders  together  the  gases  will  combine  to  form  solid  am- 
monium chloride. 

Aqua  Ammonia.  The  solubility  of  ammonia  in  water  has 
already  been  mentioned.  It  dissolves,  in  part,  without  chemical 
change,  as  is  shown  by  the  strong  odor  of  the  solution  due  to 
the  escape  of  the  gas,  but  it  partly  combines  with  the  water, 
which  easily  separates  into  H+  and  OH~,  forming  ammonium 
hydroxide  : 

H 


\OH 

This  ionizes  to  form  ammonium,  NH4+,  and  hydroxide,  OH~, 
ions,  but  the  ionization  is  small  in  comparison  with  that  of 
strong  bases.  A  tenth  normal  solution  of  sodium  hydroxide  is 
ionized  to  the  extent  of  about  84  per  cent,  while  a  tenth  normal 
solution  of  ammonium  hydroxide  (or  ammonia)  shows  only  1.3 
per  cent  of  ionization,  if  we  assume  that  all  of  the  ammonia  in 
the  solution  has  combined  with  water  to  form  ammonium 


204  A  TEXTBOOK  OF  CHEMISTRY 

hydroxide.  A  normal  solution,  on  the  same  basis,  shows  an 
ionization  of  only  0.3  to  0.4  per  cent.  How  may  the  presence 
of  hydroxide  ions  in  a  solution  of  ammonium  hydroxide  be 
demonstrated  ? 

It  is  probable,  however,  that  a  large  part  of  the  ammonia 
exists  as  such  in  the  solution.1  In  other  words,  the  solution 
contains  ammonia,  NH3,  as  well  as  ammonium  hydroxide, 
NH4OH,  and  ammonium,  NH4+,  and  hydroxide,  OH~,  ions. 
The  practical  effect,  that  ammonium  hydroxide  is  a  weak  base 
because  its  solution  contains  few  hydroxide  ions,  is  the  same 
whether  we  suppose  this  to  be  because  ammonium  hydroxide 
is  only  slightly  ionized  or  whether  it  is  because  the  ammonium 
hydroxide  is  largely  dissociated  into  ammonia  and  water. 

Ice  Machines.  Ammonia  may  be  readily  condensed  to  a 
liquid  either  by  pressure  (4.19  atmospheres  at  0°)  or  by  cold 
(—  33°  at  760  mm.).  The  heat  of  vaporization  of  the  liquid 
is  330  calories  per  kilogram  (at  —  33°).  This  high  value  is 
intimately  connected  with  the  low  molecular  weight  and  also 
with  the  fact  that  liquid  ammonia  is,  like  water,  a  highly  "  asso- 
ciated "  liquid,  that  is,  consists  of  polymerized  molecules  such 
as  (NH3)2  or  (NH3)3.  Much  heat  is  absorbed  in  the  vaporiza- 
tion both  because  of  the  large  volume  of  the  vapor  in  proportion 
to  its  weight  and  because  the  polymerized  molecules  must  be 
broken  up.  The  high  heat  of  vaporization  is  utilized  in  ice 
machines.  The  principle  of  one  form  of  these  machines  is  illus- 
trated in  the  diagram  (Fig.  71).  Liquefied  anhydrous  ammonia 
is  allowed  to  evaporate  in  the  coil  A  and  the  escaping  gas  is  com- 
pressed by  the  pump  B  and  condenses  to  a  liquid  in  the  coil 
C.  From  this  coil  the  liquid  ammonia  is  returned  through  the 

1  Moore,  J.,  Chem.  Soc.  91,  1382  (1907),  calculates  on  the  basis 
of  the  partition  of  ammonia  between  water  and  chloroform  at  dif- 
ferent temperatures  that  from  30  to  40  per  cent  of  the  ammonia 
exists  as  ammonium  hydroxide,  NH4OH,  at  20°.  This  does  not 
seem  to  be  consistent,  however,  with  the  relative  ionization  constants 
of  trimethyl  amine,  (CH3)3N,  and  tetramethyl  ammonium  hydroxide, 
(CH3)4NOH,  which  indicate  strongly  that  the  former  forms  only  a 
small  amount  of  trimethyl  ammonium  hydroxide,  (CH3)3NHOH,  in 
aqueous  solutions.  Further  experimental  evidence  on  this  ques- 
tion seems  highly  desirable. 


DERIVATIVES  OF  AMMONIA 


205 


regulating  valve  D  to  the  coil  A.  The  coil  C  is  surrounded  by 
cold  water  to  absorb  the  heat  evolved  as  the  ammonia  condenses. 
The  coil  A  is  surrounded  with  brine,  either  a  solution  of  salt 
or  of  calcium  chloride.  The  cold  brine  may  be  circulated  by 
means  of  pumps  through  coils  of  pipe  in  refrigerator  rooms,  or 
cases  of  distilled  water  may  be  immersed  in  the  brine,  to  be 
frozen. 

In  another  form  of  machine,  which  was  in  earlier  use,  the 
pressure  to  condense  the  ammonia  was  obtained  by  heating  a 


Fig.  71 

concentrated  aqueous  solution.  Afterwards  the  weakened 
solution  was  cooled,  and  as  it  reabsorbed  the  gas  the  evapora- 
tion of  the  liquefied  gas  caused  the  refrigeration. 

*  Derivatives  of  Ammonia.  One  or  more  atoms  of  hydrogen 
in  ammonia  may  be  replaced  by  a  metal,  giving  such  compounds 
as  sodium  amide,  NaNH2,  or  by  radicals,  especially  by  organic 
radicals,  giving  such  compounds  as  methyl  amine,  CHaNH^, 
phenyl  amine  or  aniline,  CeHsNH^,  acetamide,  C2H3ONH2,  and 


phthalimide,  C6H 


NH.    If  the  group  re'placing  hydrogen 


is  a  hydrocarbon  radical,  as  methyl,  CH3,  or  phenyl,  C6H6,  the 
compound  is  called  an  amine.  The  amines  combine  with  acids, 
as  ammonia  does,  to  form  such  salts  as  methyl  ammonium 
chloride,  CH3NH3C1  (or  CH3NH2.HC1)  and  phenyl  ammonium 
chloride,  C6H5NH3C1.  For  this  reason  the  amines  are  often 
called  organic  bases,  but  just  as  ammonia  is  a  base  only  when  it 


206  A  TEXTBOOK  OF  CHEMISTRY 

has  combined  with  water  to  form  ammonium  hydroxide,  NH4OH, 
the  amines  are  true  bases  only  when  combined  with  water.  As 
ammonium  hydroxide  dissociates  to  ammonia  and  water  even 
in  solution,  such  hydroxides  as  methyl  ammonium  hydroxide, 
CH3NH3OH,  dissociate  readily  into  water  and  the  original 
amine  and  can  exist  as  pure  compounds  only  at  very  low  tempera- 
tures, if  at  all.  All  four  of  the  hydrogen  atoms  in  the  ammo- 
nium group,  NH4,  of  ammonium  hydroxide  may  be  replaced  by 
hydrocarbon  radicals,  however,  and  some  compounds  formed  in 
this  manner  no  longer  dissociate  when  their  solutions  are  evapo- 
rated. Thus  a  white  crystalline  mass,  doubtless  consisting  of 
tetramethyl  ammonium  hydroxide,  (CH3)4NOH,  separates  on 
evaporating  a  solution  of  this  compound.  The  preparation  of 
this  and  several  other  similar  compounds  has  been  one  of  the 
reasons  for  believing  that  solutions  of  ammonia  in  water  contain 
ammonium  hydroxide,  NH4OH. 

If  the  group  replacing  a  hydrogen  atom  in  ammonia  is  an 
acid  radical,  the  compound  is  called  an  amide.  Thus  the  com- 
pound containing  the  acetyl  group,  C2H3O,  is  called  acetamide, 
C2H3ONH2.  In  aqueous  solutions  the  amides  are  usually  am- 
photeric;  that  is,  they  have  both  very  weak  acid  and  very  weak 
basic  properties.  When  two  hydrogen  atoms  of  ammonia  have 
been  replaced  by  a  bivalent  acid  radical,  the  compound  is 

/co\ 

called   an   imide,   as   phthalimide,    CeH^         />NH.      In    the 

XXX 
imides    the    hydrogen    can    be    replaced    by    metals    forming 

C°\ 


well-defined  salts,  as  potassium  phthalimide,  C6H4^         ">NK 

or  C6H4<(   ^>N. 

The  Electron  Theory.     A  comparison  of  the  compounds  of 
nitrogen  on  the  basis  of  the  electron  theory  (p.  181)  is  suggestive. 

In   nitric   acid,  H— O— N^     ,  the  nitrogen  atom  gives  five 


ELECTRON  THEORY  207 

electrons  to  the  oxygen  atoms,   forming  a  compound  which 

i^0= 

readily  ionizes  to  hydrogen,   H+,  and  nitrate,    O  —  N+>^      » 

+XT 

ions  because  of  the  strongly  positive  nitrogen  atom,  which  holds 
the  oxygen  of  the  hydroxyl  firmly  but  repels  its  hydrogen.     In 

/H 

ammonia,   H— N<^     ,  the  nitrogen  atom  receives  three  elec- 


trons  from  the  hydrogen  atoms,  becoming  negative.  It  can 
receive  a  fourth  electron  from  another  hydrogen  atom  only  in 
case  it  also  gives  up  one  electron  to  the  oxygen  of  a  hydroxyl 
group  or  to  chlorine  or  some  other  negative  atom  or  group.  In 

H\          /H+ 
the  ammonium  hydroxide,  H  —  ?  N+\  ,  which  results, 

H+/         X~O-H+ 

in  the  first  case,  the  negative  nitrogen  atom  no  longer  holds  the 
negative  oxygen  of  the  hydroxyl  group  strongly,  and  so  the  com- 
pound may  ionize  to  ammonium,  NHU+,  and  hydroxide,  OH~,  ions. 
If  one  of  the  hydrogen  atoms  of  ammonia  is  replaced  by  a  nega- 
tive radical,  as  acetyl,  CzH-sQ,  giving  the  compound  acetamide, 
C2H3Ov  ,  the  nitrogen  atom  no  longer  takes  up  hydrogen 

H^N 

W 

and  hydroxyl  readily  to  form  a  base,  or  the  elements  of  an 
acid  to  form  a  salt.  This  seems  to  be  because  the  presence  of 
the  negative  acetyl  group,  C2H3O,  so  far  reduces  the  positive 
character  of  the  group  C2H3(X  H  that  it  cannot  form 


W 

the  positive  ion  of  a  salt. 

*  Solutions  in  Liquid  Ammonia.  Anhydrous  ammonia  may 
be  condensed  to  a  liquid  which  boils  at  —  33.5°.  This  liquid 
ammonia  dissolves  many  substances,  and  the  conductivity  of 
the  solutions  indicates  that  some  of  these  ionize  in  the  ammonia 
as  acids,  bases  and  salts  ionize  in  solutions  in  water.  While  a 
very  large  portion  of  our  study  of  chemistry  deals  with  reactions 


208  A  TEXTBOOK  OF  CHEMISTRY 

in  aqueous  solutions,  there  are  closely  parallel  phenomena  in 
ammoniacal  solutions.  In  such  solutions  we  may  consider 
ammonia  as  consisting  of  H  and  NH2  just  as  we  think  of  water  as 
consisting  of  H  and  OH.  As  in  aqueous  solutions  derivatives 
of  water  ionize  to  form  hydrogen,  H+,  or  hydroxide,  OH~,  ions 
according  to  the  nature  of  the  radical,  so  in  ammonia,  com- 
pounds which  are  derivatives  of  ammonia  may  ionize  to  form  hy- 
drogen, H+,  or  amide,  NH2~,  ions.  Thus  acetamide, 
C2H3ONH2,  ionizes  to  C2H3ONH-  and  H+  in  solution 
in  ammonia  and  is  to  be  considered  as  an  acid  in  such  a 
solution.  Sodium  amide,  NaNH2,  on  the  other  hand, 
ionizes  to  sodium,  Na+,  and  amide,  NH2~,  ions  and  is  to 
be  considered  as  a  base.  Curiously  enough  the  latter  will 
cause  phenolphthalein  to  turn  red  in  the  ammonia  solu- 
tion just  as  hydroxide  ions  cause  it  to  turn  red  in  aqueous 
solutions.  Neutralization  in  such  a  solution  must  con- 
sist in  the  union  of  hydrogen,  H+,  and  amide,  NH2~,  ions 
to  form  ammonia.  See  Franklin  and  his  coworkers, 
Am.  Chem.  J.  20,  820  and  826 ;  21,  8 ;  23,  277 ;  28,  83 ; 
47,  285;  J.  Am.  Chem.  Soc.,  26,  499;  27,  192,  820. 

The  Volumetric  Composition  of  Ammonia.  The  ratio 
between  nitrogen  and  hydrogen  in  ammonia  may  be 
demonstrated  by  filling  a  tube,  Fig.  72,  with  chlorine 
gas,  allowing  a  small  amount  of  concentrated  aqua  am- 
monia to  enter  it  and  following  this  with  some  dilute 
sulfuric  acid.  On  allowing  water  to  enter  the  tube  till 
the  gas  remaining  in  it  is  at  atmospheric  pressure,  the 
Fig.  72  tube  will  be  found  to  be  one  third  full  of  nitrogen. 
Under  the  conditions  of  the  experiment  the  tube  full  of 
chlorine,  which  we  know  is  capable  of  combining  with  its  own 
volume  of  hydrogen,  takes  this  amount  of  hydrogen  from  the 
ammonia  and  liberates  the  equivalent  amount  of  nitrogen.  In 
other  words,  one  volume  of  nitrogen  is  combined  with  three 
volumes  of  hydrogen. 

As  the  gram  molecular  volume  of  ammonia  weighs  17  grams, 
the  complete  reaction  is  : 


COMPOSITION   OF  AMMONIA 


209 


NH3 

NH3 

The  primary  reaction  between  ammonia  and  chlorine  is : 

3  NH3  +  6  C12  =  NC13  +  N2  +  9  HC1 
This  is  followed  by  the  reaction  : 

NC13  +  NH3  =  N2  +  3  HC1 
% 

which  is  favored  in  an  acid  solution.1  The  final  result  is  the 
same  as  though  the  simple  reaction 

2  NH3  +  3  C12  =  6  HC1  +  N2 

took  place.  As  in  this  reaction  three  molecules  of  chlorine,  C12, 
liberate  one  molecule  of  nitrogen,  N2,  it  is  evident  that  in  accord- 
ance with  Avogadro's  law  a  tube  full  of  chlorine  should  liberate 
from  ammonia  one  third  of  a  tube  full  of  nitrogen. 

The  composition  of  ammonia  by  volume  may  also  be  shown 
by  another,  quite  different,  experiment.  If  a  small  amount  of 
ammonia  gas  is  introduced  into  the  apparatus  shown  in  Fig.  73, 
and  electric  sparks  are  passed  between  the  platinum  wires,  which 
pass  through  the  walls  of  the  tube,  the  ammonia  will  be  decom- 
posed into  hydrogen  and  nitrogen.  The  reaction  is  reversible, 
but  with  the  equilibrium  very  far  toward  the  side  of  decomposi- 
tion at  the  temperature  of  the  electric  discharge : 

2  NH3  ±£  N2  +  3  H2 

As  two  molecules  of  ammonia  give  one  molecule  of  nitrogen 
and  three  molecules  of  hydrogen,  the  volume  of  the  gas  would 
be  doubled  when  the  decomposition  was  complete. 

1  J.  Am.  Chem.  Soc.  23,  460. 


210 


A  TEXTBOOK  OF  CHEMISTRY 


On  the  other  hand,  if  a  mixture  of  one  volume  of  nitrogen  with 
three  volumes  of  hydrogen  is  placed  in  the  same  apparatus  and 
a  little  dilute  sulfuric  acid  is  introduced,  on  passing  electric 
sparks  through  the  mixture  as  before,  the  nitrogen  and  hydrogen 
will  slowly  combine,  and,  as  the  ammonia 
formed  will  be  absorbed  by  the  sulfuric  acid, 
the  combination  may  be  carried  to  comple- 
tion in  spite  of  the  unfavorable  character  of 
the  equilibrium. 

Nitric  Acid.  The  two  compounds,  ammonia 
and  nitric  acid,  are  to  be  considered  as  the 
fundamental  ones  for  nitrogen.  All  other  com- 
pounds of  the  element  tend  to  return  to  one  or 
the  other  of  these,  or  their  salts,  or  else  to  de- 
compose with  the  liberation  of  free  nitrogen. 
And  all  compounds  of  nitrogen  prepared  in  the 
laboratory,  except  those  derived  from  organic 
materials,  are  prepared  directly  or  indirectly 
from  ammonia  or  nitric  acid.  It  seems  best, 
therefore,  to  speak  of  nitric  acid  next,  though 
such  an  order  of  treatment  is  practical  rather 
than  logical. 

The  formation  of  nitrates  in  the  soil  by  the 
action  of  nitrifying  organisms  has  been  referred 
to.  The  present  commercial  source  for  nitrates 
is  almost  exclusively  the  sodium  nitrate,  NaNOs, 
or  Chile  saltpeter,  found  in  enormous  beds  in  Chile,  in  South 
America.  From  this  nitric  acid  is  prepared  by  a  process  similar 
to  that  for  the  preparation  of  hydrochloric  acid.  Nitric  acid  is  a 
stronger  acid  than  sulfuric,  but  if  sodium  nitrate  is  mixed  with  sul- 
furic acid  and  the  mixture  heated,  the  equilibrium  of  the  reaction, 

NaN03  +  H2SO4  ^±  NaHSO4  +  HNO3 

is  displaced  to  the  right  as  the  nitric  acid  distills  away  from  the 
mixture,  the  boiling  point  of  nitric  acid  being  much  lower  than 
that  of  the  sulfuric  acid.  Nitric  acid  is  not,  however,  very  stable, 


Fig.  73 


NITRIC  ACID 


211 


and  part  of  it  decomposes,  forming  oxygen,  water  and  nitrogen 
peroxide,  NO2,  when  the  distillation  is  under  atmospheric  pres- 
sure. To  avoid  this  the  operation  is  sometimes  carried  out  at  a 
lower  temperature  by  reducing  the  pressure. 

Pure  nitric  acid  is  a  colorless  liquid,  which  boils  at  86°  and 
has  a  specific  gravity  of  1.52.  The  addition  of  water  causes  a 
rise  in  the  boiling  point,  the  acid  of  maximum  boiling  point, 
120°,  containing  66  to  70  per  cent  of  pure  acid. 

Nitric  acid  is  a  strong  acid,  the  tenth  normal  solution  being 
ionized  to  the  extent  of  92  per  cent,  while  a  tenth  normal  solution 
of  hydrochloric  acid  is  91  per  cent  ionized. 

*  Hydrates  of  Nitric  Acid.  The  addition  of  increasing  amounts 
of  nitric  acid  to  water  lowers  the  freezing  point  till  an  acid  con- 
taining 32.8  per  cent  of  nitric  acid  freezes  at  —  43°.  Further 
addition  of  nitric  acid  causes  the  freezing  point  to  rise  and  fall 


°C0 

S     10 

*S 

\ 

1 
«     ~2° 

E 

r,.            or* 

\ 

\ 

\ 

/ 

\ 

TEMPERATURE  Ol 

i  A  i  I  c 

0  0  0  0  C 

\ 

j 

\ 

V 

\ 

/ 

\ 

1 

)        10      20       30       40      50      60      70      80      90     100% 
PER  CENT.  OF  NITRIC  ACID.              HNO3 

.     Fig.  74 

alternately  as  shown  in  Fig.  74,  pure  nitric  acid  freezing  at 
—  41.2°.  The  two  maximum  freezing  points  shown  in  the  figure 
correspond  to  acids  containing  53.84  and  77.77  per  cent  of 


212  A  TEXTBOOK  OF  CHEMISTRY 

nitric  acid  respectively.  In  accordance  with  the  principle  that 
a  pure  liquid  freezes  at  a  higher  temperature  than  when  it  con- 
tains some  dissolved  substance,  these  melting  points  indicate 
definite  compounds.  An  acid  containing  53.84  per  cent  of  nitric 
acid  corresponds  to  the  formula  HNO3.3  H2O  and  one  contain- 
ing 77.77  per  cent,  to  the  formula  HNO3.H2O.  The  freezing 
point  curve  demonstrates  very  clearly  the  existence  of  these  two 
hydrates.1 

Chemical  Properties  of  Nitric  Acid.  As  an  acid,  nitric  acid 
has  the  usual  properties,  forming  salts  as  a  monobasic  acid  with 
practically  all  metals.  These  may  be  prepared  by  the  action  of 
the  acid  on  the  metal  or  on  a  hydroxide,  oxide  or  carbonate  of 
the  metal : 

NaOH  +  HN03  =  NaNO3  +  HOH 
Ca(OH)2  +  2  HNO3  =  Ca(NO3)2  +  2  H2O 
ZnO  +  2  HNO3  =  Zn(NO3)2  +  H2O 

The  most  important  special  properties  of  nitric  acid  depend 
on  the  ease  with  which  it  gives  up  oxygen  to  a  great  variety  of 
substances.  It  is  a  powerful  oxidizing  agent,  the  concentrated 
or  the  anhydrous  acid,  HNO3,  showing  this  property  in  a  more 
marked  degree  than  the  dilute  acid.  Indeed  the  pure  acid  de- 
composes spontaneously  into  water,  oxygen  and  nitrogen  perox- 
ide, NO2,  on  distillation  or  on  exposure  to  light : 

2  HNO3  =  2  NO2+  O  +  H2O 

If  pure  nitric  acid  is  boiled  in  a  test  tube  having  a  plug  of  wool 
or  feathers  in  its  mouth,  the  latter  will  catch  fire  and  burn. 
Ignited  charcoal  will  also  continue  to  burn  beneath  the  surface 
of  the  liquid. 

When  nitric  acid  acts  upon  a  metal,  hydrogen  is  rarely,  if  ever, 
liberated.  Instead  of  this  the  nitric  acid  is  reduced,  either  by 
the  metal  directly,  forming  an  oxide  of  the  metal,  or  by  the  hy- 
drogen displaced  by  the  metal.  Such  hydrogen  is  often  called 

1  Kiister  u.  Kremann,  Z.  anorg.  Chem.  41,  I  (1904). 


AQUA  REGIA  .  213 


nascent"1  hydrogen.     The  two  ways  of  explaining  the  action 
may  be  illustrated  as  follows  : 

3  Cu  +  2  HNO3  =  (3  CuO)  +  H2O  +  2  NO 

(3  CuO)  +  6  HN03  =  3  Cu(NO3)2  +  3  H2O 

3  Cu  +  8  HNO3  =  3  Cu(NO3)2  +  4  H2O  +  2  NO 
Or  3  Cu  +  6  HN03  =  3  Cu(NO3)2  +(6  H) 

(6  H)  +  2  HNO3  =  4  H20  +  2  NO 

3  Cu  +  8  HNO3  =3  Cu(NO3)2  +  2  NO  +  4  H2O 

The  final  reaction  is  the  same  whichever  explanation  of  the 
mechanism  of  the  reaction  is  adopted. 

Nitric  acid  may  be  reduced  less  or  more  by  other  metals,  or  by 
copper  when  the  concentration  of  the  nitric  acid  is  different, 
giving  the  whole  series  of  oxides,  NO2,  N2O3,  NO  and  N2O,  and 
even  ammonia,  NH3.  In  general  a  concentrated  acid  gives  the 
higher  oxides,  while  a  more  dilute  acid  or  a  metal  which  has  a 
high  heat  of  oxidation  gives  the  lower  oxides  or  ammonia.  If 
nitric  acid  is  added  to  a  test  tube  containing  zinc  and  sulf  uric  acid, 
the  evolution  of  hydrogen  may  nearly  cease  and  ammonium 
sulf  ate  will  be  formed.  The  addition  of  an  excess  of  sodium 
hydroxide  to  the  solution  will  then  liberate  ammonia,  which  may 
be  recognized  by  its  odor  and  effect  on  litmus  paper.  What  is 
the  series  of  reactions  involved  in  the  experiment  ? 

Aqua  Regia.  Neither  nitric  acid  or  hydrochloric  acid  alone 
will  dissolve  the  so-called  noble  metals,  gold  or  platinum,  but  a 
mixture  of  the  two  will  dissolve  them  readily.  Such  a  mixture 
is  called  aqua  regia 2  because  of  this  property.  The  nitric  acid 
acts  upon  hydrochloric  acid  in  the  same  manner  as  other  oxidiz- 
ing agents,  liberating  chlorine,  and  this  attacks  the  gold  or  plati- 
num, forming  soluble  chlorides. 

Because  it  contains  chlorine  and  oxides  of  nitrogen,  aqua  regia 

1  Nascent  means  in  the  state  of  being  born.  The  theory  is  that 
hydrogen  atoms  when  first  liberated  are  more  active  than  the  same 
atoms  when  they  have  combined  with  others  to  form  hydrogen 
molecules,  H2.  The  activity  is  dependent  largely  on  the  metal  used, 
however,  metals  which  dissolve  with  a  large  evolution  of  heat  giving 
the  most  active  hydrogen.  2  Royal  water. 


214  A   TEXTBOOK  OF  CHEMISTRY 

is  a  powerful  oxidizing  agent  and  is  frequently  used  for  that  pur- 
pose, especially  to  oxidize  the  sulfur  of  sulfides  to  sulfuric  acid, 
as  in  the  determination  of  sulfur  in  iron. 

Nitrosyl  chloride,  NOC1,  a  volatile  compound  with  a  very 
disagreeable  odor,  is  also  formed  in  the  mixture  of  nitric  and 
hydrochloric  acids.  This  is  hydrolyzed  by  water  in  the  same 
manner  as  other  nonmetallic  chlorides  : 

NOC1  +  HOH  =  NOOH  (or  HNO2)  +  HC1 

Nitrosyl  chloride  is  also  a  strong  oxidizing  agent. 

Oxides  of  Nitrogen.  There  are  six  oxides  of  nitrogen,  but  as 
two  of  these,  nitrogen  peroxide,  NO2,  and  nitrogen  tetroxide, 
N2O4,  have  the  same  percentage  composition  and  change  each 
into  the  other  on  merely  changing  the  temperature,  they  are 
frequently  spoken  of  as  a  single  substance  and  the  name  nitrogen 
peroxide  is  applied  to  both.  The  oxides  are : 

Nitrous  oxide  N2O 

Nitric  oxide  NO 

Nitrous  anhydride  N2O3 

Nitrogen  dioxide  NO2 

Nitrogen  tetroxide  N2O4 

Nitric  anhydride  N2O5 

Nitrous  oxide,  N2O,  is  most  easily  prepared  by  heating  am- 
monium nitrate;  nitric  oxide,  NO,  by  the  reduction  of  nitric 
acid  with  metallic  copper ;  nitrous  anhydride,  N2O3,  by  reducing 
nitric  acid  with  arsenious  oxide,  As2Os ;  nitrogen  dioxide,  NO2, 
by  heating  lead  nitrate,  Pb(NO3)2;  and  nitric  anhydride,  N2O5, 
by  dehydrating  nitric  acid  with  phosphoric  anhydride.  Nitric 
oxide  is  also  formed  by  the  direct  union  of  the  elements  in  an 
electric  arc,  and  it  unites  with  oxygen  spontaneously  to  form 
nitrous  anhydride,  N2O3,  and  nitrogen  dioxide,  NO2. 

Nitrous  Oxide,  N2O.  When  ammonium  nitrate,  NH4NO3, 
is  heated,  the  hydrogen  of  one  part  of  the  molecule  combines 
with  oxygen  from  another  part  to  form  water,  while  the  two 
nitrogen  atoms  remain  combined  with  the  other  oxygen  atom : 

NH4NO3  =  N2O  +  2  H2O 


NITRIC  OXIDE  215 

The  reaction  is  exothermic,  that  is,  it  takes  place  with  evolu- 
tion of  heat  and  is  liable  to  become  explosive  if  the  temperature 
is  raised  too  high  or  if  too  large  a  quantity  of  the  salt  is  heated 
at  once. 

Nitrous  oxide  is  a  colorless  gas  with  a  sweetish  odor  and  taste. 
Water  at  20°  absorbs  about  two  thirds  of  its  volume  of  the  gas. 
It  supports  combustion.  A  glowing  splinter  will  inflame  in 
the  gas  somewhat  as  it  does  in  oxygen,  and  phosphorus  burns  in 
it  with  an  intense  light.  It  will  not  support  life. 

When  inhaled,  nitrous  oxide  sometimes  causes  hysterical  laugh- 
ing, and  it  is  called  for  that  reason  laughing  gas.  In  larger 
amounts  it  produces  insensibility  and  is  used  for  this  purpose 
in  minor  surgical  operations,  especially  for  the  extraction  of 
teeth. 

The  structure  of  nitrous  oxide  is  probably  that  represented  by 

N 

the  formula,  II  }O.     This  accounts  best  for  the  ease  with  which 
W 

it  gives  up  oxygen  and  reverts  to  free  nitrogen.  It  is  the  only 
oxide  of  nitrogen  in  which  two  atoms  of  nitrogen  are  supposed 
to  be  directly  united. 

Nitric  Oxide,  NO,  is  formed  when  copper  dissolves  in  dilute 
nitric  acid  of  specific  gravity  1.2.  If  a  stronger  acid  is  used,  some 
nitrogen  dioxide,  NO2,  will  be  formed ;  while  if  the  acid  is  much 
more  dilute,  nitrous  oxide,  N2O,  and  nitrogen,  N2,  will  be  formed 
along  with  the  nitric  oxide.  The  mechanism  of  the  reaction  has 
already  been  discussed. 

Nitric  oxide,  NO,  is  also  formed  by  the  direct  union  of  nitrogen 
and  oxygen  in  an  electric  arc.  For  instance,'  if  electric  sparks 
from  an  induction  coil  are  passed  for  some  time  between  terminals 
in  a  large  globe,  the  air  in  the  globe  will  gradually  become  reddish 
brown  in  color  from  the  formation  of  nitric  oxide,  which  combines 
with  more  oxygen  to  form  nitrogen  dioxide,  NO2.  There. is  evi- 
dence that  the  combination  is  caused  by  the  high  temperature 
of  the  arc  and  not  by  the  electricity  as  such. 

The  combination  of  nitrogen  and  oxygen  is  an  endothermic 


216 


A  TEXTBOOK  OF  CHEMISTRY 


reaction  ;  that  is,  heat  is  absorbed  as  it  proceeds.     It  is  also  a 
reversible  reaction  : 


with  the  equilibrium  very  far  to  the  left,  so  far,  indeed,  that  the 
amount  of  nitric  oxide  formed  from  the  elements  is  very  small, 
even  at  high  temperatures.  The  per  cent  of  nitric  oxide  formed 
in  air  when  the  reaction  comes  to  equilibrium  is  as  follows  :  1 


ABSOLUTE 
TEMPERATURE            .. 

PER  CENT  OF  NO 
CALCULATED 

OBSERVED 

1811° 

0.35 

0.37 

2195° 

0.98 

0.97 

2675° 

2.37 

2.23 

3200° 

4.43 

About  5. 

With  the  equilibrium  so  far  on  the  side  toward  its  decomposi- 
tion, it  seems  at  first  thought  that  nitric  oxide  ought  not  to  exist 
at  all  at  ordinary  temperatures,  and  it  could  not  except  for  the 
fact  that  the  speed  of  the  formation  or  decomposition  is  very 
slow.  Thus  it  has  been  shown  that  for  the  formation  of  half  of 
the  amount  corresponding  to  a  condition  of  equilibrium,  80  years 
would  be  required  at  a  temperature  of  725°  and  one  and  a  fourth 
days  at  1225°.  At  1825°  it  takes  only  5  seconds. 

These  facts  are  important  in  determining  the  heat  conditions 
for  the  preparation  of  nitric  oxide  from  the  air  as  the  first  step  in 
manufacturing  nitrates.  Evidently  the  highest  possible  tem- 
perature must  be  secured,  and  when  the  nitric  oxide  has  been 
formed  at  that  temperature,  it  must  be  cooled  as  quickly  as 
possible  through  the  temperatures  at  which  formation  and  de- 
composition are  both  rapid  and  where  the  point  of  equilibrium 
lies  farther  toward  the  side  of  decomposition. 

Nitric  oxide  is  a  colorless  gas  which  may  be  condensed  to  a 
liquid  that  boils  at  -153.6°  and  freezes  at  -167°.  Is  the  gas 
heavier  or  lighter  than  air  ? 

1  Nernst,  Z.  anorgf.  Chem.  49,  213. 


NITRIC  OXIDE  217 

Nitric  oxide,  quite  unlike  nitrous  oxide,  will  extinguish  a 
glowing  coal  or  a  candle  which  is  thrust  into  it,  or  even  a  piece 
of  phosphorus  which  is  barely  ignited.  If  a  little  hotter,  phos- 
phorus burns  brilliantly  in  the  gas,  and  a  mixture  of  the  vapor 
of  carbon  bisulfide  with  the  gas  will  burn  with  a  bright  blue  flash, 
which  is  rich  in  the  light  rays  that  affect  a  photographic  plate. 

Nitric  oxide  combines  directly  with  oxygen  to  form  nitrous 
anhydride,  N2O3,  nitrogen  dioxide,  NO2,  or  nitrogen  tetroxide, 
N2O4,  or  a  mixture  of  the  three,  according  to  the  temperature, 
and  the  relative  volumes  of  the  gases.  If  a  strong  base,  such 
as  potassium  hydroxide,  KOH,  is  present  when  nitric  oxide  and 
oxygen  are  brought  together,  a  nitrite,  KNO2,  is  formed,  even 
though  the  oxygen  is  in  excess.  In  what  proportion  by  volume 
must  nitric  oxide  and  oxygen  be  brought  together  to  form  nitrous 
anhydride  ?  In  what  proportion  to  form  nitrogen  dioxide  ? 

The  structural  formula  of  nitric  oxide  is  usually  written 
N  =  O,  representing  the  nitrogen  as  bivalent,  a  valence  which  is 
very  unusual  for  the  element  and  which  it  is  not  known  to  have 
in  any  other  compound.  This  unusual  structure  seems  to  be 
closely  connected  with  the  tendency  of  nitric  oxide  to  combine 
with  more  oxygen. 

*  Nitric  oxide  is  formed  by  the  reduction  of  nitric  acid  by 
ferrous  sulfate,  FeSO4,  the  latter  being  oxidized  to  ferric  sulfate, 
Fe2(SO4)3,  if  sulfuric  acid  is  present : 


2  HNO3  =  2  NO  +  (3  O)  +  H2O 
2  FeSO4  +  H2SO4  +  (O)  =  Fe2(SO4)3  +  H2O 


Combining  the  equations, 

6  FeSO4  +  3  H2SO4  +  2  HNO3  =  3  Fe(SO4)3  +  3  H2O  +  2  NO 

The  nitric  oxide  is  absorbed  by  a  solution  of  ferrous  sulfate 
with  the  formation  of  an  unstable  compound,  FeSO4.NO 
(see  Manchot  and  Zechentmayer,  Ann.  350,  368),  which  gives 
to  the  solution  a  dark  brown  or  black  color  and  which  is  often 
used  for  the  qualitative  detection  of  nitric  acid.  The  reduction 


218  A  TEXTBOOK  OF  CHEMISTRY 

of  nitric  acid  to  nitric  oxide  by  ferrous  chloride,  FeCl2,  and  hydro- 
chloric  acid  is  also  used  for  its  quantitative  determination. 

Nitrous  anhydride,  N2O3,  may  be  prepared  by  the  union  of 
oxygen  with  nitric  oxide,  by  the  action  of  a  dilute  acid  on  sodium 
nitrite,  NaNO2,  or  by  the  reduction  of  nitric  acid,  especially  by 
arsenious  oxide,  As2O3  : 

4  NO  +  O2  =  2  N2O3 

2  NaNO2  +  H2SO4  =  Na2SO4  +  2  HNO2 

Nitrous 
Acid 

2  HNO2  =  N2O3  +  H2O 
2  HNO3  +  As2O3  +  2  H2O  =  2  H3AsO4  +  N2O, 

Arsenic 
Acid 

At  ordinary  temperatures  a  gas  having  the  composition  of 
nitrous  anhydride  has  a  density  which  indicates  that  the  com- 
pound dissociates  into  nitric  oxide,  NO,  nitrogen  dioxide,  NC>2, 
and  nitrogen  tetroxide,  N2O4  : 


2  NO2  ^±  N2O4     (See  below.) 

At  a  low  temperature  the  gases  recombine,  in  part,  and  con- 
dense to  a  dark  blue  or  green  liquid,  but  even  at  —  90°  there  is 
still  some  dissociation  (Ramsay). 

Nitrous  Acid.  If  the  mixture  of  gases  spoken  of  in  the  last 
paragraph  is  dissolved  in  cold  water,  nitrous  acid,  HNO2,  is 
formed  : 

NO  +  NO2  +  H2O  ^±  2  HN02 

The  acid  is  very  unstable  and  can  exist  only  in  dilute  solutions. 
Salts  of  nitrous  acid  are  most  easily  prepared  by  reducing  a  ni- 
trate, as,  for  instance,  sodium  nitrate,  with  lead  or  copper  : 

NaNO3  +  Pb  =  NaNO2  +  PbO 

Sodium 
Nitrite 


NITROGEN  DIOXIDE  219 

Nitrogen  Dioxide,  NO2,  and  Nitrogen  Tetroxide,  N2O4. 
When  nitric  oxide  is  mixed  with  air  or  oxygen,  the  colorless  gas 
changes  to  a  reddish  brown  color  and  is  converted,  at  ordinary 
temperatures,  into  a  mixture  of  nitrogen  dioxide,  NO2,  and 
nitrogen  tetroxide,  N2O4.  The  molecular  weight  of  nitrogen 
dioxide  is  46,  that  of  nitrogen  tetroxide  is  92.  As  the  gram 
molecular  volume  of  the  gas  weighs  about  80  grams  at  10°, 
69  grams  at  64°  and  46  grams  at  150°,  it  is  evident  that  at  the 
lower  temperature  it  consists  chiefly  of  nitrogen  tetroxide,  N2O4, 
but  that  this  dissociates  into  nitrogen  dioxide,  NO2,  as  the 
temperature  rises  : 


At  still  higher  temperatures  nitrogen  dioxide  dissociates  into 
nitric  oxide,  NO,  and  oxygen. 


The  mixture  of  the  two  gases  may  be  easily  condensed  to  a 
reddish  brown  liquid  which  boils  at  25°.  At  lower  temperatures 
the  liquid  becomes  lighter  colored  and  solidifies  to  colorless  crys- 
tals at  —  10.5°.  From  this  it  would  seem  that  nitrogen  tetrox- 
ide is  colorless,  while  the  dioxide  is  colored. 

The  structure  of  the  two  compounds  is  probably  : 


. 

and     ^N—  O—  N=O 

CT 


The  tetroxide,      \N — O — N=O,    may    be    considered    as 
partly  nitrous  anhydride,  O=N — O — N=O,  and  partly  nitric 


anhydride,       ^N — O — N^     •     In  accordance  with  this  view 

CT 

it  gives   both   nitric  and   nitrous    acid  when  dissolved  in  cold 
water : 


220  A  TEXTBOOK  OF  CHEMISTRY 


V 


/i 


N— O— N=O  ,O 

# 


=  H— O— NC    +H— O— N=0 
H— O— H  ^O 

Nitric  Acid  Nitrous  Acid 


In  warm  water  nitrogen  dioxide  forms  nitric  acid  and  nitric 
oxide  :  ,/>O 

NO2  +  HOH  =  H—  O—  Nf     +  (H) 


2(H)+  NO2  =  NO  +  H2O 
or  3  NO2  +  H2O  =  2  HNO3  +  NO 

As  the  nitric  oxide  on  coming  to  the  air  immediately  combines 
with  oxygen  to  form  the  dioxide,  these  reactions  furnish  a  means 
for  converting  the  nitric  oxide  obtained  from  atmospheric  nitro- 
gen in  the  electric  arc  into  nitric  acid. 

Nitrogen  Pentoxide  is  formed  when  nitric  acid  is  mixed  with 
phosphorus  pentoxide  : 

2HNO3  +  P2O6     =     2HPO3     +     N2O6 

Metaphosphoric 
Acid 

It  forms  colorless  crystals  which  melt  at  29.5°.  It  decomposes 
easily  into  nitrogen  dioxide  and  oxygen.  With  water  it  gives 
nitric  acid. 

Other  Compounds  of  Nitrogen.  The  halogens  are  univalent 
toward  hydrogen  and  form  only  one  compound,  each,  with  that 
element.  Hydrochloric  acid,  HC1,  is  typical  for  the  group. 

Oxygen,  which  is  bivalent,  forms  two  compounds,  with  hydro- 
gen H—  O—  H  and  H—  O—  O—  H. 

Nitrogen,  which  is  trivalent,  forms  at  least  five  compounds 
with  hydrogen.  ,^ 

Ammonia,  NHa,         or        Nr—  H 

\H 

H\ 


Hydrazine,1  N2H4,        or 

W  H 

1  Nitrogen  is  called  "azote"  (deprived  of  life)  in  French,  and 
many  names  of  nitrogen  compounds  are  derived  from  the  root  "  az  " 
or  "azo." 


HYDROXYLAMINE  221 

Hydronitric  acid,  N3H,         or         H— 


v 
H\ 

,  ,         H-^N 

W 


N 
N 


Ammonium  trinitride,    N4EU,    or  H-N  —  N^ 

XN 

H 


H  AH 

>N—  N^- 


Hydrazine  trinitride,      N5H5,    or       >N—  N^-  --  N 

W  XH 

We  shall  find  later  that  carbon,  which  is  quadrivalent,  forms 
many  hundred  of  compounds  with  hydrogen. 

Nitrogen  also  forms  a  third  oxygen  acid,  hyponitrous  acid, 
H2N2O2,  and  a  derivative  of  ammonia,  hydroxylamine,  NH2OH. 

*  Hyponitrous    Acid,    H2N2O2.     From    its    formula,   nitrous 
oxide,  N2O,  might  be  looked  upon  as  the  anhydride  of  hyponi- 
trous acid,  but  it  will  not  combine  with  water  to  form  the  acid 
nor  will  it  combine  with  bases  to  form  salts  of  the  acid.     The 
method  of  preparation  which  indicates  most  clearly  the  structure 
of  the  acid  is  by  the  interaction  of  hydroxylamine  and  nitrous 
acid:  ---------  . 

'  H 

H—  O—  N<^    +0=N—  OH  =  H—  O—  N=N—  O—  H  +  H2O 
>H         ! 


The  pure  acid  can  be  obtained  in  crystalline  form,  but  is  very 
explosive.  The  lowering  of  the  freezing  point  of  an  aqueous 
solution  of  the  acid  (see  p.  112)  shows  that  the  formula  is  H2N2O2, 
and  not  the  simpler  formula,  HNO.  In  solutions  of  either 
alkalies  or  acids  it  decomposes  slowly  into  water  and  nitrous 
oxide,  N2O. 

*  Hydroxylamine,  NH2OH,  may  be  prepared  by  the  electro- 
lytic reduction  of  nitric  acid,  using  an  amalgamated  lead  cathode 
and  in  the  presence  of  sulfuric  acid,  which  combines  with  the 
hydroxylamine  to  form  hydroxylammonium  sulfate  : l 
1  Tafel,  Z.  anorg.  Chem.  81,  289. 


222  A  TEXTBOOK  OF  CHEMISTRY 

,0  K 

H—  O—  NC     +  6  H  =  H—  O—  N<      +  2  H2O 


H 
2  NH2OH  +  H2SO4  =   (NH3OH)2SO4 

Hydroxylammonium 
Sulfate 

Pure  hydroxylamine  may  be  obtained  in  white  scales  or 
needles  which  melt  at  33°.  It  dissolves  easily  in  water  and  the 
solution  seems  to  contain  some  hydroxylammonium  hydroxide, 

/OK 
NH3^        ,  which  is  a  much  weaker  base  than  ammonium  hy- 

X)H 

droxide,  NH4OH.  Hydroxylamine  is  much  used  in  organic 
chemistry  for  the  preparation  of  derivatives  of  aldehydes  and 
ketones,  called  oximes. 

*  Hydrazine,  N2H4.  By  means  of  a  series  of  reactions  a  com- 
pound called  bisdiazoacetic  acid  may  be  prepared.  When  this 
is  hydrolyzed  by  warming  with  hydrochloric  acid,  it  gives  hydra- 
zine  hydrochloride  and  oxalic  acid. 


X          \ 
H02C—  CH<  >CHC02H  +  2  HC1  +  4  H2O 

XN  =  NX 

Bisdiazoacetic  Acid 

CO2H 
=  2H2N—  NH2HCl  +  2    | 

Hydrazine          CO2H 
Hydrochloride    Oxalic  Acid 

Pure  hydrazine  is  a  colorless  liquid,  which  solidifies  at  a 
low  temperature  and  melts  at  1.4°.  In  solution  it  forms 
a  base,  which  seems  to  be  H2N  —  NH3OH,  similar  to  ammo- 
nium hydroxide,  NH4OH,  and  hydroxylammonium  hydroxide, 
NH3OH.OH.  With  acids  it  forms  salts  by  direct  addition, 
taking  up  either  one  or  two  molecules  of  a  monobasic  acid, 
the  chlorides  being  H2N—  NH3C1  and  C1H3N—  NH3C1.  The 
second  molecule  of  the  acid  is  held  only  feebly,  however.  In 
aqueous  solution  hydrazine  gives  a  tenfold  weaker  base  than 
ammonium  hydroxide. 


HYDRONITRIC  ACID  223 


Hydrazine  and  its  derivatives,  especially  phenylhydrazine, 
C6H6NHNH2,  are  much  used  in  organic  chemistry  for  the 
preparation  of  derivatives  of  aldehydes,  ketones  and  sugars. 


/N 

*  Hydronitric  Acid  or  Azoimide,  HN^  || .     By  passing  am- 

\XT 


monia  over  heated  metallic  sodium  a  compound  called  sod- 
amide,  NaNH2,  may  be  prepared.  When  nitrous  oxide,  N2O, 
is  passed  over  this  at  a  temperature  of  190°,  sodium  trinitride, 
NaN3,  is  formed  : 

/H         /N  /N 

Na—  N<     +O<  ||   =  Na—  N<  ||  +H2O 
XH         XN  XN 

By  dissolving  the  sodium  trinitride  in  water,  adding  sulfuric  acid, 
and  distilling,  the  hydronitric  acid  passes  over  with  some  water, 
from  which  it  can  be  separated  partly  by  redistillation  and  finally 
by  treating  with  calcium  chloride  : 

2  NaN3  +  H2SO4  =  Na2SO4  +  2  HN3 

The  pure  acid  freezes  at  —  80  and  boils  at  +  37°.  It  is  very 
explosive,  as  are  also  many  of  its  salts,  especially  the  salts  with 
silver,  AgN3,  and  copper,  Cu(N3)2.  The  silver  and  copper  salts 
are  very  difficultly  soluble  in  water.  In  aqueous  solution  the  free 
acid  is  only  slightly  ionized,  its  strength  in  this  respect  resem- 
bling that  of  acetic  acid.  It  is,  however,  a  much  stronger  acid 
than  hydrogen  sulfide,  H2S,  or  carbonic  acid,  H2CO3. 

Iodine  Trinitride.  By  treating  silver  trinitride  with  iodine 
the  silver  may  be  replaced  by  iodine  : 


, 

Ag-N< 
X 


N 

Agl  +  I- 

N 


Iodine  trinitride  is  a  colorless  solid  with  a  penetrating  odor, 
resembling  that  of  cyanogen  iodide.  It  is  hydrolyzed  by  water 
to  hydronitric  and  hypoiodous  acids  : 


224  A  TEXTBOOK   OF  CHEMISTRY 


I— N<  ||  +  H.OH  =  H— N<  ||  +1— O— H 
XN  NN 

This  reaction  is  remarkable  because  the  iodine  conducts  itself 
as  the  positive  part  of  the  molecule,  combining  with  the  negative 
hydroxyl,  while  the  hydrazoic  group,  N3,  combines  with  the  posi- 
tive hydrogen.  Probably  for  the  same  reason  silver  trinitride 
will  not  react  with  iodine  trinitride  to  give  the  molecule  N6, 

N — N<^  ||     ,  because  both  the  silver  and  iodine  are  positive 

in  the  two  compounds. 

Nitrogen  Trichloride.  When  a  dilute  solution  of  ammonia 
reacts  with  chlorine,  one  fourth  of  the  chlorine  combines  with 
nitrogen  to  form  nitrogen  trichloride : 

3  NH3  +  6  Cla  =  NC13  +  9  HC1  +  N2 
9  HC1  +  9  NH3  =  9  NH4C1 

Nitrogen  trichloride  is  also  formed  when  hypochlorous  acid 
acts  on  ammonium  chloride  : 

H  Cl 

N^-H  +  3  HO.C1  ^±  3  H.OH  +  N^-C1 

XH  XC1 

The  last  reaction  seems  to  be  reversible,  and  these  reactions  indi- 
cate that  the  chlorine  of  nitrogen  trichloride  is  positive,  just  as 
the  iodine  of  iodine  trinitride  is.  Because  of  this,  one  molecule 
of  nitrogen  trichloride  is  equivalent  to  six  atoms  (or  three  mole- 
cules, 3  C12)  of  available  chlorine  in  oxidizing  power : 

2  NC13  +  3  As2O3  +  15  H2O  =  6  H3AsO4  +  2  NH3  +  6  HC1 
6  C12  +  3  As2O3  +  15  H2O  =  6  H3AsO4  +  12  HC1 

Nitrogen  chloride  is  a  volatile  oil,  which  is  very  explosive  and 
dangerous.  Dulong,  who  discovered  it,  lost  an  eye  and  three 
fingers  while  working  with  it,  and  both  Faraday  and  Davy  were 


ENDOTHERMIC  COMPOUNDS  225 

injured  while  experimenting  with  the  substance.  It  is  soluble 
in  benzene  and  may  be  handled  more  safely  in  such  a  solution. 

Nitro  Nitrogen  Trichloride.  Some  evidence  has  been  obtained 
recently,  which  points  to  the  existence  of  a  nitro  nitrogen  tri- 
chloride in  which  the  nitrogen  is  positive  and  the  chlorine  nega- 
tive. See  J.  Am.  Chem.  Soc.  35,  767  (1913). 

*  Nitrogen  Iodide,  N2H3l3,  is  formed  when  a  strong  solution  of 
ammonia  is  poured  over  powdered  iodine.  When  dried  in  the 
air  it  forms  a  black  powder  which  explodes  violently  at  the 
lightest  touch. 

Endothermic  Compounds.  Troost  has  given  the  heat  of  forma- 
tion of  nitrogen  chloride  from  its  elements  as  —38,477  calories. 
The  explosive  character  of  nitrogen  chloride  and  nitrogen  iodide 
is  evidently  closely  connected  with  the  fact  that  they  are  en- 
dothermic  compounds,  that  is,  compounds  which  are  formed  from 
their  elements  with  absorption  of  heat  and  which,  conversely, 
decompose  into  their  elements  with  evolution  of  heat.  The 
absorption  of  heat  in  their  formation  indicates  that  the  atoms 
have  little  affinity  for  each  other  in  the  compounds  and  so  may 
easily  separate,  and  the  heat  generated  by  the  decomposition, 
when  it  once  begins,  raises  the  mixture  rapidly  to  a  higher 
temperature  and  this  in  turn  hastens  the  reaction  till  it  becomes 
explosive.  The  explosion  is  further  caused,  of  course,  by  the 
formation  of  a  large  volume  of  gas  at  a  high  temperature  from 
a  small  volume  of  a  liquid  or  solid. 

EXERCISES 

1.  The  specific  gravity  of  mercury  at  0°  is  13.6.     What  is  the  weight 
of  air  above  one  square  centimeter  of  the  earth's  surface  at  sea  level  ? 

2.  Sketch  an  apparatus  suitable  for  the  preparation  of  nitrogen  by 
the  use  of  hydrogen  and  copper  oxide. 

3.  What  is  the  weight  of  a  gram  molecular  volume  of  air  if  it  contains 
21  per  cent  of  oxygen,  0.9  per  cent  of  argon  and  78.1  per  cent  of  nitro- 
gen ?     What  is  the  weight  of  one  liter  ? 

4.  If  a  gram  molecular  volume  of  the  mixture  of  nitrogen  peroxide 
and  nitrogen  tetroxide  weighs  69  grams  at  64°,  what  is  the  per  cent  of 
each  compound  present  ? 


226  A  TEXTBOOK  OF  CHEMISTRY 

5.  How  many  grams  of  nitric  acid  will  be  required  to  dissolve  5 
grams  of  copper  ?   How  many  liters  of  nitric  oxide  will  be  formed  ? 

6.  What  weight  of  ammonia  will  be  contained  in  one  liter  of  aqua 
ammonia  having  a  specific  gravity  of  0.90  and  containing  28  per  cent  of 
ammonia  ?     How  much  water  ?     How  many  grams  of  ammonia  would 
a  liter  of  water  take  up  in  forming  a  solution  of  specific  gravity  0.90  ? 
How  many  liters  of  the  gas  ? 

7.  How  much  salt,  NaCl,  will  be  required  to  furnish  the  hydrochloric 
acid  necessary  to  neutralize  17  grams  of  ammonia?     How  much  sul- 
f uric  acid  ? 

8.  If  an  isomeric  nitrogen  trichloride,  NC13,  could  be  prepared  in 
which  the  chlorine  atoms  were  negative,  what  would  be  the  products 
of  its  hydrolysis? 


CHAPTER  XIII 
THE   ATMOSPHERE.     NOBLE   GASES 

Determination  of  Oxgyen.  The  first  determination  of  the 
per  cent  of  oxygen  in  the  atmosphere  was  made  by  Lavoisier, 
who  heated  a  measured  quantity  of  air  with  mercury  as  long  as 
the  oxygen  continued  to  be  absorbed  (p.  19). 
He  determined  roughly  the  contraction  which 
occurred  and  also  showed  that  by  heating  the 
oxide  of  mercury  formed  he  obtained  a  volume 
of  oxygen  closely  agreeing  with  this  contrac- 
tion. His  results  showed  that  approximately 
one  fifth  of  the  volume  of  air  is  oxygen.  The 
determination  may  be  made  more  accurately 
by  measuring  a  volume  of  air  in  a  eudiometer, 
exposing  it  to  the  action  of  a  stick  of  phos- 
phorus, which  is  slowly  oxidized  to  phospho- 
rous acid,  H3PO3,  and  measuring  the  volume 
of  the  gases  which  remain  (Fig.  75).  Or  a 
measured  volume  of  air  may  be  mixed  with  a 
little  more  than  two  fifths  of  its  volume  of 
hydrogen  and  the  mixture  exploded.  One 
third  of  the  contraction  in  volume  (or,  more 

accurately  ,  see  p.  68)  will  be  the  volume 

o.OOZL 

of  the  oxygen  in  the  air  taken. 

Composition    of    Air.     Very    many    careful 
analyses  of  air  have  shown  that  when  samples          Fig.  75 
are  taken  out  of  doors  under  usual  conditions 
the  composition  of  dry  air  agrees  very  closely  with  the  fol- 
lowing : 

227 


228 


A  TEXTBOOK  OF  CHEMISTRY 


BY  VOLUME 

BY  WEIGHT 

Oxvsren 

20  95  per  cent 

23.15  per  cent 

Carbon  dioxide  

0.03  per  cent 

0.05  per  cent 

Argon 

0  94  per  cent 

1.3    per  cent 

Nitrogen   

78.08  per  cent 

75.5    per  cent 

100 

100 

The  percentage  of  oxygen  varies  slightly,  but  is  very  rarely 
less  than  20.9  or  more  than  21.0  per  cent,  out  of  doors. 

Air  is  a  Mixture.  Although  the  composition  of  air  is  nearly 
constant,  it  is  believed  that  the  oxygen  and  nitrogen  which  it 
contains  are  merely  mixed  together  and  not  chemically  com- 
bined, for  the  following  reasons  : 

1.  No  heat  is  generated  when  oxygen,  nitrogen,  argon  and 
carbon  dioxide  are  mixed,  and  yet  the  mixture  has  all  of  the 
properties  of  air. 

2.  Liquefied  air  when  it  boils  gives  at  first  a  gas  richer  in 
nitrogen  and  at  last  nearly  pure  oxygen.     A  compound  when 
it  boils  away  without  decomposition  has  the  same  composition 
from  first  to  last. 

3.  When  water  is  shaken  with  air,  it  dissolves  the  oxygen  and 
nitrogen  in  proportion  to  the  solubility  and  partial  pressure 
of  each,  and  as  oxygen  is  more  soluble  than  nitrogen,  the  mixture 
of  gases  obtained  by  boiling  the  water  contains  about  35  per  cent 
of  oxygen,  while  air  contains  only  21  per  cent.     This  is  the 
conduct  to  be  expected  of  a  mixture  rather  than  that  of  a  com- 
pound. 

4.  A  gram  molecular  volume  of  air  weighs  28.95  grams.     The 
simplest  formula  which  would  give  a  composition  approximating 
that  of  air  is  N4O.     The  gram  molecular  volume  of  a  compound 
of  this  formula  would  weigh  72  grams. 

On  the  other  hand,  a  liter  of  air  weighs  1.2928  grams,  which  is 
almost  exactly  what  it  should  weigh  if  it  is  a  mixture  of  gases 
in  the  proportion  which  has  been  given,  as  will  be  seen  from  the 
following  calculation : 


THE  ATMOSPHERE  229 

WEIGHT  or  ONE 

LITER] 

Oxygen  1.429    X  0.2095  =  0.2994 

Carbon  dioxide  1.9768  X  0.0003  =  0.0006 

Argon  1.7828  X  0.0094  =  0.0168 

Nitrogen  1.2507  X  0.7808  =  0.9765 

1.2933 

Carbon  Dioxide  in  the  Air.  The  carbon  dioxide  in  the  air 
comes  from  four  principal  sources  :  1.  From  the  breath  of  men 
and  animals,  the  carbon  of  food  which  is  eaten  being  mostly 
converted  into  carbon  dioxide  by  the  oxidation  processes  which 
take  place  in  the  body.  2.  From  the  burning  of  compounds 
containing  carbon,  such  as  wood,  coal,  oil  or  natural  gas. 
3.  From  the  decay  of  animal  and  vegetable  substances  under 
the  influence  of  bacteria.  4.  From  volcanoes  and  other  sub- 
terranean sources. 

While  the  amount  of  carbon  dioxide  in  the  air  is  very  small 
in  comparison  with  the  amounts  of  oxygen  and  nitrogen,  the  total 
amount  is  very  great.  Thus  it  is  estimated  that  1,300,000,000 
tons  of  coal  are  burned  annually,  and  this  gives  nearly  three 
times  its  weight  of  carbon  dioxide,  but  this  would  increase  the 
amount  in  the  air  by  only  one  six-hundredth  part.1  The  carbon 
dioxide  of  the  air  is  also  a  very  important  factor  in  the  economy 
of  nature,  as  it  furnishes  practically  all  of  the  carbon  for  the 
growth  of  plants.  In  a  sense  it  is  the  constituent  of  the  atmos- 
phere which  is  most  vitally  important  for  the  life  of  plants, 
as  oxygen  is  the  constituent  necessary  for  the  life  of  animals. 
In  another  sense  it  may  be  said  to  be  even  more  important,  since 
it  furnishes  the  most  important  constituent  for  the  growth  of 
plants,  while  oxygen  furnishes  to  animals  only  the  means  with 
which  to  consume  food  which  is  secured  from  some  other  source. 
In  utilizing  carbon  dioxide  plants  decompose  it  and  exhale  oxygen 
to  the  air.  In  this  way  the  growth  of  plants  prevents  the  accu- 
mulation of  carbon  dioxide  in  the  atmosphere.  The  process  of 

1  A.  Krogh,  quoted  byj  F.  W.  Clarke  in  Data  of  Geochemistry, 
p.  42.     See  also  Science,  1911,  p.  757. 


230  A  TEXTBOOK  OF  CHEMISTRY 

reducing  carbon  dioxide  to  the  compounds  of  carbon  synthe- 
sized by  the  plant  is  endothermic,  of  course,  and  the  necessary 
energy  is  furnished  by  the  sunlight.  When  we  burn  coal  in  our 
furnaces,  we  make  use  of  the  energy  of  sunlight  which  was  stored 
by  growing  plants  millions  of  years  ago.  It  has  even  been  sug- 
gested that  all  of  the  oxygen  of  the  air  came  originally  from  car- 
bon dioxide  through  this  process  of  plant  growth  (Lord  Kelvin). 

Besides  the  equilibrium  maintained  by  the  balance  between 
the  evolution  of  carbon  dioxide  from  the  sources  named  and  the 
absorption  of  the  gas  by  growing  plants,  the  ocean  plays  a  very 
important  part  in  maintaining  a  constant  amount  in  the  air 
during  long  periods.  As  carbon  dioxide  dissolves  in  water  in 
proportion  to  the  partial  pressure  of  the  gas  (p.  165),  any  increase 
in  the  amount  of  carbon  dioxide  in  the  air  would  be  followed 
very  quickly  by  an  increase  in  the  amount  in  the  ocean,  while 
any  decrease  would  be  replaced  from  the  storehouse  in  the  ocean. 
Since  the  amount  of  "  free  "  carbon  dioxide  in  the  ocean  to  a 
depth  of  5  kilometers  is  nearly  fifty  times  1  the  amount  in  the  air 
above  it,  and  about  three  fourths  of  the  earth's  surface  is  covered 
by  the  ocean,  the  importance  of  the  store  contained  in  the  ocean 
is  obvious. 

Ventilation.  Shortly  after  the  discovery  of  the  composition 
of  the  air  by  Priestly  and  Lavoisier,  a  method  was  devised  for 
analyzing  air  by  mixing  it  with  nitric  oxide,  to  combine  with  the 
oxygen,  and  then  absorbing  the  nitrogen  peroxide  formed  by 
means  of  a  solution  of  potassium  hydroxide.  It  will  be  readily 
understood  that  such  a  method  requires  very  great  care  to  secure 
accurate  results  and  the  early  determinations  led  the  observers 
to  think  that  there  was  a  considerable  fluctuation  in  the  amount 
of  oxygen  present  and  that  this  fluctuation  caused  the  difference 
between  good  and  bad  air.  But  Cavendish  was  able  to  use  even 
the  nitric  oxide  method  so  accurately  that  he  very  soon  showed 
that  the  variation  in  the  composition  of  the  atmosphere  must 
be  between  very  narrow  limits,  and  this  result  has  been  confirmed 

1  Calculated  on  the  basis  of  45  milligrams  per  liter  of  sea  water. 


VENTILATION  231 

by  later  observers.  It  was  then  discovered  that  carbon  dioxide 
in  mines  and  in  wells  or  caves  frequently  killed  persons  exposed 
to  its  action,  and  for  many  years  it  was  supposed  that  this  gas 
acts  as  a  positive  poison  and  is  the  chief  cause  of  danger  in  poorly 
ventilated  rooms.  This  fallacy  and  also  the  opinion  that  carbon 
dioxide  will  accumulate  near  the  floor  of  a  room  because  the  gas 
is  one  and  a  half  times  as  heavy  as  air  were  spread  so  widely  in 
semipopular  literature  and  became  so  firmly  fixed  in  the  minds 
of  many  people  that  it  has  proved  very  difficult  to  correct  these 
errors.  It  has  been  shown  that  the  amount  of  carbon  dioxide 
present  in  the  air  of  even  badly  ventilated  rooms  is  practically 
never  great  enough  to  cause  any  injury  to  human  beings.  It  has 
been  found  very  difficult  to  demonstrate  clearly  just  what  sub- 
stances cause  the  ill  effects  which  follow  from  poor  ventilation, 
and  some  recent  authorities  have  spoken  doubtfully  of  the 
standards  for  ventilation  which  have  been  proposed.  There 
seems  to  be  little  doubt,  however,  that  lack  of  ventilation  in 
factories,  offices  and  dwellings  is  a  frequent  cause  of  disease.  It 
is  also  very  well  established  that  abundance  of  fresh  air  secured 
by  life  out  of  doors,  both  by  night  and  day,  combined  with  a 
nourishing  diet,  furnish  the  best  hope  of  recovery  from  incipient 
tuberculosis. 

While  exhaled  carbon  dioxide  is  not  in  itself  harmful,  it  fur- 
nishes the  best  means  of  determining  whether  a  room  occupied 
by  people  is  properly  ventilated  or  not.  The  amount  of  the 
gas  should  not  exceed  0.07  per  cent  by  volume.  To  secure 
this  amount  of  ventilation  55,000  liters  or  2000  cubic  feet  of 
fresh  air  will  be  required  each  hour  for  each  person  in  a  room.1 

Moisture.  Natural  air  always  contains  a  certain  amount  of 
water  vapor,  but  this  is  subject  to  very  great  variations,  depen- 
dent on  the  temperature  and  the  conditions  to  which  the  air  has 
been  subjected.  The  pressure  of  the  water  vapor  can  never 
much  exceed  the  normal  vapor  pressure  of  water  for  the  given 

1  Roscoe  and  Schorlenmer,  Treatise  on  Chemistry,  I,  589. 
Another  authority  recommends  85,000  liters  per  hour.  Stewart, 
Manual  of  Physiology,  p.  244. 


232  A  TEXTBOOK  OF  CHEMISTRY 

temperature,  and,  indeed,  can  only  exceed  that  when  in  a  state 
of  unstable  equilibrium  such  that  the  introduction  of  suitable 
nuclei  to  form  points  of  condensation  will  at  once  cause  the  for- 
mation of  a  cloud.  When  the  pressure  of  the  water  vapor  in  the  air 
corresponds  to  the  normal  pressure  of  water  vapor  for  the  given 
temperature,  the  air  is  said  to  be  saturated,  but  such  a  condition 
does  not  usually  obtain  close  to  the  earth's  surface.  At  a  height 
of  a  few  hundred  or  thousand  feet,  however,  owing  partly  to  the 
mixing  of  warm,  nearly  saturated  air  with  colder  air  currents, 
partly  to  the  lowering  of  the  temperature,  which  results  from 
the  adiabatic  cooling  of  air  as  it  expands  in  rising,  saturation 
and  condensation  to  clouds  and  rain  take  place. 

The  amount  of  moisture  in  air  may  be  determined :  1.  By 
aspirating  a  known  volume  through  weighed  bulbs  containing 
concentrated  sulfuric  acid.  2.  By  determining  the  dew  point, 
that  is,  the  temperature  at  which  the  air  will  deposit  moisture  on 
a  cooled,  polished,  metallic  surface  or  the  temperature  at  which 
moisture  will  just  disappear  from  such  a  surface.  3.  By  com- 
paring the  temperature  of  the  air  with  the  temperature  of  a 
thermometer  whose  bulb  is  covered  with  moist  cotton  over  which 
air  is  blown.  Tables  have  been  prepared  giving  the  humidity 
corresponding  to  the  difference  observed. 

In  general  the  humidity  in  rooms  which  are  heated  is  too  low 
for  healthfulness  and  should  be  supplemented  by  artificial  means. 
In  many  factories,  especially  in  those  for  spinning  and  weaving, 
the  degree  of  humidity  is  of  vital  importance  to  the  success  of  the 
operations. 

Liquid  Air.  Critical  Temperature.  After  it  had  been  shown 
early  in  the  nineteenth  century  that  such  gases  as  chlorine, 
ammonia,  carbon  dioxide,  and  sulfur  dioxide  could  be  liquefied, 
many  attempts  were  made  to  liquefy  oxygen  and  nitrogen,  or 
air.  These  gases  were  subjected  to  pressures  of  several  hundreds 
of  atmospheres,  but  it  was  always  found  that  the  gas  continued 
to  fill  completely  and  uniformly  any  space  left  to  it,  while  if 
liquefied  it  should  have  separated  into  a  liquid  and  a  gaseous 
portion.  Finally,  in  1869,  Andrews  showed  that  carbon  dioxide, 


CRITICAL  TEMPERATURE  233 


which  can  be  liquefied  under  a  pressure  of  38.5  atmospheres  at 
0°,  or  71  atmospheres  at  30°,  cannot  be  liquefied,  even  under 
pressures  very  much  greater  than  this,  at  temperatures  above  31°. 
If  a  thick-walled,  sealed  glass  tube  containing  liquid  carbon  diox- 
ide is  warmed  gently,  at  a  temperature  of  31.35°  the  liquid  in  the 
lower  part  of  the  tube  will  suddenly  disappear  and  the  gas  will 
now  fill  the  tube  uniformly.  The  pressure  may  be  increased  or 
decreased,  but  as  long  as  the  temperature  is  above  31.35°  no  pres- 
sure either  high  or  low  can  be  found  at  which  the  carbon  dioxide 
will  separate  into  a  liquid  phase  and  a  vapor  phase.  Below  this 
temperature  carbon  dioxide  will  be  partly  liquid  and  partly  gas, 
provided  the  pressure  is  equal  to  the  vapor  pressure  of  the  liquid 
at  the  given  temperature  and  the  volume  filled  by  the  substance 
is  large  enough  to  allow  a  part  to  assume  the  vapor  phase.  The 
temperature  above  which  a  gas  cannot  be  liquefied  is  called  the 
critical  temperature. 

Andrews'  experiment  made  it  seem  very  probable  that  the  fail- 
ures to  liquefy  air  were  due  to  the  fact  that  the  critical  tempera- 
tures of  oxygen  and  nitrogen  are  much  below  ordinary  tempera- 
ture. Following  this  suggestion,  Cailletet  in  Paris  and  Pictet 
in  Geneva  (1877),  working  independently,  both  succeeded  in 
liquefying  oxygen  by  the  use  of  cold  and  pressure  combined  with 
the  cooling  effect  produced  by  the  expansion  of  the  highly  com- 
pressed gas. 

Some  years  later  it  was  shown  by  Joule  and  Thomson  (Lord 
Kelvin)  that  a  moderately  compressed  gas  scarcely  changes  its 
temperature  on  expanding  into  a  vacuum  —  for  instance,  if  air 
compressed  to  20  atmospheres  is  allowed  to  expand  into  a  vacu- 
ous receptacle,  both  receptacles  being  surroun'ded  by  water,  the 
temperature  scarcely  changes,  though  for  all  gases  except  hydro- 
gen and  helium,  there  is  a  slight  cooling  effect.  This  cooling  effect 
increases  for  higher  pressures,  or  when  a  gas  is  so  far  compressed 
that  it  no  longer  obeys  Boyle's  law  (p.  35).  It  would  seem  that 
the  attraction  between  the  molecules  of  the  gas  has  a  greater 
effect  as  the  molecules  are  brought  closer  together,  causing  the 
gas  to  contract  more  than  it  should  in  accordance  with  the  law. 


234 


A  TEXTBOOK  OF  CHEMISTRY 


When  the  gas  expands  from  such  a  condition  work  must  be  done 
in  overcoming  this  attraction  between  the  molecules,  and  the 
expansion  is  accompanied  by  a  cooling  effect.  On  the  basis  of 
these  facts  Linde,  Hampson  and  others  have  devised  machines 
by  means  of  which  air  can  be  readily  liquefied  in  large  quantities. 
_  In  these  machines  air  is  compressed  to 

150-200  atmospheres  and  is  then  allowed 
to  expand  to  atmospheric  pressure  in  such 
a  manner  that  the  expanded  and  cooled 
air  passes  back  over  the  tube  in  which 
the  air  is  expanding.  In  the  Hampson 
machine  the  air  expands  through  a  copper 
tube  of  about  three  millimeters  in  internal 
diameter  and  one  hundred  and  thirty 
meters  in  length.  This  is  wound  in  a 
spiral  to  secure  compactness  and  the  ex- 
panded air  is  compelled  to  follow  the 
course  of  the  spiral  backwards,  Fig.  76. 
By  these  machines  a  portion  of  the  air  is 
soon  cooled  to  the  point  of  liquefaction 
and  the  liquid  air  collects  in  a  receptacle 
placed  beneath  the  end  of  the  spiral. 

The  carbon  dioxide  must  be  removed 
from  the  air  which  is  to  be  liquefied,  by 
passing  it  through  a  large  apparatus  filled 
with  slaked  lime,  and  the  moisture  must 
also  be  removed  by  calcium  chloride  or 
some  drying  agent,  as  otherwise  these 
would  condense  in  solid  form  and  stop  up 

the  tube  through  which  the  air  expands.  For  the  liquefaction  of 
hydrogen  the  compressed  gas  must  be  cooled  by  liquid  air,  as 
it  is  only  at  low  temperatures  that  hydrogen  depart^  sufficiently 
from  Boyle's  law  so  that  it  can  be  liquefied  by  this  method. 

Liquid  nitrogen  boils  at  —194°,  liquid  oxygen  at  —182.5°. 
Liquid  air  will  contain,  therefore,  a  larger  proportion  of  oxygen 
than  ordinary  air,  and  by  a  sort  of  fractional  distillation  it  is  easy 


Fig.  76 


LIQUID  AIR.    ARGON  235 

to  obtain  from  it  a  gas  which  contains  from  75  to  95  per  cent  of 

oxygen.     Such  a  gas  is  used  for  medicinal  (e.g.  in  pneumonia) 

and  some  technical  purposes.     The  method 

is  also  used  to  obtain  nearly  pure  nitrogen 

and  is  now  the  most  important  industrial 

method  for  the  preparation  of  both  oxygen 

and  nitrogen. 

For  experimental  purposes  liquid  air  is 
kept  in  Dewar  flasks  (Fig.  77),  double- 
walled  flasks  having  the  space  between  the 
two  walls  evacuated  to  prevent  loss  of  heat 
by  convection  currents.  The  inner  bulb  is 
often  silvered  to  cause  it  to  reflect  radiant  Fig.  77 

heat  which  reaches  it  from  outside. 

Argon,  A,  39.88.  In  1785  Cavendish  described  an  experiment 
in  which  he  mixed  air  with  an  excess  of  oxygen,  passed  electric 
sparks  through  the  mixture,  and  absorbed  the  oxides  of  nitrogen 
formed  by  a  solution  of  potassium  hydroxide.  He  then  absorbed 
the  rest  of  the  oxygen  by  means  of  "  liver  of  sulfur  "  and  re- 
ported that  the  gas  remaining  unabsorbed  was  not  more  than 
Tffr  of  the  original  volume  of  the  air.  The  real  significance  of 
this  remarkable  experiment  was  not  understood  for  more  than 
a  century. 

During  the  eighties  and  nineties  of  the  last  century  Lord 
Rayleigh  spent  a  great  deal  of  time  in  determining  very  accu- 
rately the  density  of  the  elementary  gases,  oxygen,  hydrogen 
and  nitrogen.  In  the  course  of  his  work  he  prepared  what  he  sup- 
posed to  be  nitrogen  by  removing  oxygen  and  all  other  known 
substances  from  the  air.  He  also  prepared  nitrogen  by  the 
decomposition  of  ammonia.  To  his  surprise  a  liter  of  the  nitro- 
gen obtained  from  the  air  weighed  about  6  milligrams  more  than 
a  liter  of  nitrogen  prepared  from  ammonia.  Lord  Rayleigh  is 
a  physicist,  and  he  called  in  the  assistance  of  a  chemist,  Sir 
William  Ramsay,  to  solve  the  problem  which  was  presented. 
Within  a  short  time,  in  1894,  the  two  succeeded  in  preparing 
argon,  partly  by  a  repetition  of  the  Cavendish  experiment  with 


236  A  TEXTBOOK  OF  CHEMISTRY 

modern  appliances,  partly  by  removing  the  nitrogen  of  the  air 
by  passing  it  over  heated  magnesium,  with  which  the  nitrogen 
combined. 

Argon  was  not  only  a  new  element,  but  it  belongs  to  a  wholly 
new  class  of  elements,  now  called  the  Zero  group  of  the  Periodic 
System,  or  the  noble  gases.  The  most  remarkable  property  of 
these  elements  is  that  none  of  them  enters  into  chemical  combina- 
tion with  other  elements  or  with  itself  —  the  valence  of  the 
group  is  zero. 

Argon  may  be  condensed  to  a  liquid,  which  freezes  at  — 188° 
and  boils  at  —186.1°.  The  gram  molecular  volume  weighs 
39.9  grams.  From  this  the  molecular  weight  is  39.88. 

Atomic  Weight  of  Argon.  Specific  Heat  of  Gases.  The  fact 
that  argon  will  not  combine  with  any  other  element  would,  of 
itself,  lead  us  to  expect  that  the  molecule  of  argon  consists  of  a 
single  atom  and  that  the  formula  of  the  gas  is  A  and  the  atomic 
weight  39.88.  Another,  wholly  independent,  line  of  evidence 
points  to  the  same  conclusion.  The  specific  heat  of  a  gas  may  be 
determined  either  while  the  volume  of  the  gas  remains  constant 
or  while  the  pressure  remains  constant.  It  is  evident  that  the 
specific  heat  must  be  greater  at  constant  pressure  than  at  con- 
stant volume  because  at  constant  pressure  the  gas  must  expand 
as  it  grows  warm  and  do  work  as  it  expands  against  the  pressure 
of  the  atmosphere.  It  can  be  shown  that,  on  the  basis  of  the 
fundamental  assumptions  of  the  kinetic  theory  of  gases,  in  any 
gas  in  which  the  energy  required  to  increase  the  temperature 
of  the  gas  is  all  used  in  increasing  the  average  velocity  of  the  mole- 
cules, the  ratio  of  the  specific  heats  must  be : 

Specific  heat  at  constant  pressure  _  1.67 
Specific  heat  at  constant  volume        1 

On  the  other  hand,  if  a  part  of  the  energy  is  used  in  causing 
the  atoms  within  these  molecules  to  vibrate  more  violently, 
the  numerator  of  the  fraction  expressing  the  ratio  between  the 
two  kinds  of  specific  heat  will  be  smaller,  since  both  kinds  of 
specific  heat  will  be  greater  and  an  addition  to  both  the  numera- 


HELIUM  237 

tor  and  denominator  of  any  fraction  causes  it  to  approach 
unity. 

The  ratio  between  the  two  specific  heats  can  be  calculated  from 
the  velocity  of  sound  in  the  gas.  The  two  kinds  of  specific  heat 
have  also  been  determined  directly  for  air  and  some  other  gases. 
It  has  been  found  that  the  ratio  of  the  specific  heats  for  mercury 
vapor,  for  argon  and  for  .some  other  gases  is  very  close  to  1.67/1 
and  it  is  believed  that  all  of  these  gases  are  monatomic.  The 
ratio  of  the  specific  heats  for  nitrogen  is  1.41/1 ;  for  carbon 
dioxide  it  is  1.305/1 ;  for  ethylene,  1.26/1 ;  and  in  general  the 
numerator  becomes  smaller  as  the  molecule  is  more  complex. 
This  seems  to  mean  that  in  gases  with  complex  molecules  a  con- 
siderable part  of  the  energy  used  in  heating  the  gas  is  absorbed 
in  doing  internal  work  in  the  molecules,  that  is,  in  causing  their 
atoms  to  vibrate  more  and  more  rapidly.  It  will  be  readily  seen 
that  this  conclusion  gives  a  simple  explanation  of  the  fact  that 
complex  molecules  are  generally  unstable  at  high  temperatures. 

Helium,  He,  3.99.  In  1868  Lockyer  observed  some  bright 
lines  in  the  spectrum  of  the  corona  of  the  sun,  which  did  not 
correspond  to  the  lines  of  any  element  then  known.  He  called 
the  element  which  gives  these  lines  helium  (from  ^Aios,  the  sun), 
and  he  had,  in  reality,  discovered  a  new  element,  which  for 
nearly  thirty  years  was  known  to  exist  only  in  the  sun,  90,000,- 
000  miles  away.  Shortly  after  the  discovery  of  argon  it  was 
recalled  that  Dr.  Hillebrand  of  the  U.  S.  Geological  Survey  had 
obtained  a  gas  from  the  mineral  uraninite.  Ramsay,  on  further 
examination  of  the  gases  obtained  from  cleveite,  a  variety  of 
uraninite,  found  in  them  not  only  a  small  amount  of  argon, 
but  also  a  gas  which  gave  the  same  spectral  lines  which  had 
been  observed  in  the  light  of  the  sun's  corona,  and  he  soon 
separated  helium  from  the  mixture.  Helium  is  only  twice  as 
heavy  as  hydrogen  and  has  the  lowest  boiling  point  of  any  known 
substance  (unless  we  call  the  electron  an  element).  It  boils 
at  —268.5°  or  at  4.5°  absolute.  Helium  has  acquired  a  very 
extraordinary  interest,  also,  from  the  discovery  that  it  is  formed 
by  the  decomposition  of  radium.  In  spite  of  this  method  of 


238  A  TEXTBOOK  OF  CHEMISTRY 

formation,  radium  cannot  be  considered  as  a  compound  of  helium, 
and  no  one  has  been  able  to  induce  helium  to  combine  with  any 
other  element. 

Helium  is  found  in  all  gases  issuing  from  the  earth.  It  is 
doubtless  derived  from  radium  and  other  radioactive  elements. 
It  has  been  suggested  that  the  reason  why  only  a  very  minute 
quantity  of  helium  is  found  in  the  atmosphere  is  because,  owing 
to  the  lightness  of  the  helium  atoms,  their  kinetic  velocity  is  such 
that  they  may  fly  away  from  the  earth  into  space. 

Neon,  Krypton,  Xenon,  and  Niton.  The  following  partial 
table  of  atomic  weights  taken  from  the  periodic  system  indicates 
that  there  should  be  three  or  four  other  elements  belonging  to 
the  same  family  as  helium  and  argon : 

He    4  Li        7 

F      19  Na    23 

Cl     35.5  A    39.9  K      39 

Br    80  Rb    85.4 

I     127  Cs    133 

A  systematic  search  for  these  elements  soon  led  Ramsay  to  the 
discovery  of  neon  (Ne  =  20.2),  krypton  (Kr  =  82.92)  and 
xenon  (X  =  130.2)  as  constituents  of  the  air,  each  of  them  pres- 
ent, however,  in  only  very  small  amounts.  Several  years  later 
it  was  shown  that  an  evanescent  element  formed  by  the  disinte- 
gration of  radium  belongs  to  this  series.  The  density  of  the  gas 
has  been  determined  only  approximately  because  of  the  minute 
quantity  which  it  is  possible  to  obtain.  From  this  determina- 
tion the  atomic  weight  is  about  222.4.  The  element  is  called 
niton.  It  disintegrates  spontaneously  and  very  rapidly,  one 
half  of  it  disappearing  in  a  little  less  than  four  days. 

EXERCISES 

1.  An  adult  eats  food  containing  about  300  grams  of  carbon  daily. 
If  this  is  exhaled  as  carbon  dioxide,  CO2,  at  a  temperature  of  37°,  how 
many  liters  of  the  gas  are  exhaled  per  hour  ? 

2.  If  a  person  breathes  20  times  per  minute,  500  cc.  of  air  being  ex- 
haled at  each  respiration  and  the  exhaled  air  contains  4  per  cent  of 
carbon  dioxide,  how  many  liters  of  the  gas  are  exhaled  per  hour  ? 


THE  ATMOSPHERE.    NOBLE  GASES  239 

3.  How  does  the  volume  of  carbon  dioxide  from  a  gas  jet  burning 
3  cubic  feet  of  gas  per  hour  and  giving  an  equal  volume  of  carbon  dioxide 
compare  with  that  exhaled  by  an  adult  ? 

4.  If  outside  air  contains  0.03  per  cent  of  carbon  dioxide,  how  often 
must  the  air  in  a  room  5  meters  square  and  3  meters  high  be  changed 
in  order  that  the  amount  of  carbon  dioxide  may  not  exceed  0.07  per  cent 
when  two  persons,  each  breathing  out  carbon  dioxide  at  the  rate  of  20 
liters  per  hour,  are  present  ? 


CHAPTER  XIV 
PHOSPHORUS 

THE  atomic  weights  of  the  nonmetallic  elements  of  the  fifth, 
sixth,  seventh  and  zero  groups  of  the  periodic  system  and  of  the 
semimetallic  elements  of  the  fifth  group  are,  in  round  numbers  : 


FIFTH  GROUP 

SIXTH  GROUP 

SEVENTH  GROUP 

ZERO  GROUP 

He  4 

N  14 

O   16 

F    19 

Ne20 

P  31 

S    32 

Cl  35.5 

A    40 

As  75" 

Se  78 

Br80 

Kr83 

Sbl20 

Te  127.6 

I    127 

Xe  130 

Bi  208 

Nt222 

Phosphorus,  P,  31.04.  Occurrence.  Phosphorus,  the  second 
element  of  the  fifth  group,  is  a  very  important  element  both  for 
vegetable  and  animal  life.  It  is  an  essential  mineral  constituent 
in  soils  for  the  growth  of  plants,  and  it  is  also  an  important  ele- 
ment in  the  protoplasm  of  the  cells  and  in  the  bones  of  animals. 
The  ash  left  when  the  organic  matter  is  burned  out  of  bones  con- 
sists very  largely  of  calcium  phosphate,  Ca3(PO4)2-  The  same 
compound  is  found  mixed  with  other  substances  in  extensive  de- 
posits of  "phosphate  rock"  in  North  and  South  Carolina, 
Georgia,  Florida  and  Tennessee.  These  deposits  are  extensively 
mined  for  use  in  applying  to  soils  which  are  deficient  in  phos- 
phorus. Phosphorus  is  also  found  in  the  mineral  apatite, 
Ca5(PO4)sF  or  CasCPO^aCl,  which  has  already  been  mentioned 
in  connection  with  fluorine.  Phosphorus  compounds  are  found 
in  almost  all  iron  ores,  lessening  their  value  when  present  in 

240 


PHOSPHORUS  241 

more  than  very  small  amounts,  because  of  the  injurious  effect  of 
the  phosphorus  on  the  iron  made  from  such  ores. 

Preparation  of  Phosphorus.  When  a.  mixture  of  sand  (sili- 
con dioxide,  SiO2),  calcium  phosphate,  Ca3(PO4)2,  and  charcoal 
or  coke,  C,  is  heated  to  a  very  high  temperature  in  an  electric 
furnace,  calcium  silicate,  CaSiOa,  phosphorus,  P4,  and  carbon 
monoxide,  CO,  are  produced  : 

2  Ca3(PO4)2  +  6  SiO2  +  10  C  =  6  CaSiO3  4-  P4  +  10  CO  * 

The  phosphorus  distills  from  the  retort  in  which  the  mixture 
is  heated  and  is  condensed  and  collected  under  water.  This 
electrical  furnace  method  for  manufacturing  phosphorus  has  dis- 
placed older,  more  complicated  methods,  in  comparatively 
recent  times. 

Allotropic  Forms  of  Phosphorus.  The  phosphorus  obtained  as 
described  is  a  waxlike  solid  which  usually  has  a  slight  yellow 
color  and  this  form  is  called  "  ordinary  "  or  "  yellow  "  phos- 
phorus. When  pure  it  melts  at  44.5°  and  can  be  readily  melted 
and  cast  into  sticks  under  water.  Its  specific  gravity  is  1.8232 
at  20°.  It  boils  at  290°.  In  the  gaseous  form  a  gram  molecular 
volume  weighs  about  124  grams,  from  which  the  formula  must 
beP4. 

Ordinary  phosphorus  glows  with  a  pale  light  when  exposed  to 
moist  air.  It  may  be  distilled  with  steam,  and  a  very  minute 
quantity  may  be  detected  in  a  dark  room  by  the  use  of  these 
properties.  The  word  phosphorescence  recalls,  of  course,  the 
luminous  quality  of  the  element. 

If  ordinary  phosphorus  is  heated  to  240°-250°  in  a  closed 
vessel,  it  is  gradually,  though  not  quite  completely,  transformed 
into  the  allo tropic  variety  called  red  phosphorus.  This  was 
formerly  called  amorphous  phosphorus,  but  it  may  be  crystallized 
from  solution  in  melted  lead.  When  pure  and  free  from  yellow 

1  In  writing  this  equation  notice  that  two  molecules  of  calcium 
phosphate,  Ca3(PO4)2,  are  required  to  give  one  molecule  of  phos- 
phorus, P4.  The  rest  of  the  equation  follows  logically  from  the 
formulas  of  the  product  formed. 


242  A  TEXTBOOK  OF  CHEMISTRY 

phosphorus  its  specific  gravity  is  2.34.  It  is  not  poisonous, 
while  yellow  phosphorous  is  very  poisonous  indeed.  Yellow 
phosphorus  dissolves  readily  in  carbon  disulfide,  red  phosphorus 
does  not.  Yellow  phosphorus  must  be  kept  away  from  the  air 
and  is  usually  kept  under  water  because  of  the  very  low  kindling 
temperature.  Red  phosphorus  takes  fire  at  a  much  higher  tem- 
perature and  may  be  kept  in  open  bottles.  When  heated  to  a 
high  temperature,  red  phosphorus  distills  and  goes  back  to  the 
yellow  form,  but  at  lower  temperatures  the  vapor  pressure  of 
red  phosphorus  is  much  lower  than  that  of  the  yellow  variety. 
The  molecular  weight  of  red  phosphorus  has  not  been  deter- 
mined. 

Matches.  The  methods  of  obtaining  fire  in  use  before  the  nine- 
teenth century  were  difficult  of  application  and  people  often  sent 
to  their  neighbors  even  at  some  distance  for  coals  rather  than  to 
take  the  trouble  of  starting  a  new  fire.  Phosphorus  was  dis- 
covered, it  is  true,  in  1669  by  Brandt,  an  alchemist  of  Hamburg, 
but  it  was  not  till  1827  that  use  was  made  of  its  low  kindling 
temperature  for  the  preparation  of  matches.  For  the  first 
matches  using  phosphorus  the  match  sticks  were  dipped  in 
melted  sulfur  and  then  in  a  mixture  of  phosphorus  and  glue 
or  some  other  adhesive  substance.  When  dry  a  slight  friction 
raises  the  phosphorus  to  its  kindling  temperature  and  this,  as  it 
burns,  sets  fire  to  the  sulfur,  which,  in  turn,  ignites  the  wood. 
In  the  later  manufacture  the  sulfur  was  replaced  by  other  com- 
bustible substances  which  do  not  give  an  objectionable  odor,  and 
the  kindling  power  of  the  phosphorus  was  reenforced  by  potas- 
sium chlorate,  red  lead  or  other  oxidizing  compounds.  Ordinary 
phosphorus  is  extremely  poisonous,  however,  and  gives  off  enough 
vapor  at  ordinary  temperatures  so  that,  unless  extraordinary 
pains  are  taken  to  ventilate  the  factories,  the  workmen  often 
suffer  from  a  very  painful  and  fatal  disease,  which  causes  necro- 
sis of  the  jaw.  Partly  for  this  reason  and  partly  to  avoid  the 
danger  of  accidental  fires,  most  European  countries  have  for- 
bidden the  sale  or  even  the  manufacture  of  matches  containing 
ordinary  phosphorus.  The  "  safety  "  matches  used  in  these 


PHOSPHINE  243 

countries  have  on  their  heads,  usually,  a  mixture  of  antimony 
trisulfide,  potassium  chlorate  and  glue,  and  they  are  ignited 
on  a  prepared  surface  of  red  phosphorus,  glue  and  a  sulfide  of 
antimony. 

In  comparatively  recent  times  it  has  been  discovered  that 
tetraphosphorus  trisulfide,  P4S3,  may  be  substituted  for  yellow 
phosphorus  in  ordinary  matches.  As  it  does  not  give  off  poison- 
ous vapors,  this  sulfide  of  phosphorus  ought  soon  to  entirely 
displace  the  ordinary  phosphorus  for  this  manufacture.  A  law 
passed  by  Congress  in  1912  will  prevent  the  further  use  of  ordi- 
nary phosphorus  for  matches  in  the  United  States. 

Phosphine,  PH3.  When  yellow  phosphorus  is  warmed  with 
a  strong  solution  of  sodium  hydroxide,  it  is  oxidized  to  sodium 
hypophosphite,  NaH^PC^.  At  the  same  time  some  of  the  hy- 
drogen of  the  water  or  of  the  sodium  hydroxide  combines  with 
more  of  the  phosphorus  to  form  phosphine,  PH3.  In  preparing 
the  gas  a  rather  small  flask  should  be  used,  and  it  is  well  to  add 
to  the  contents  of  the  flask,  before  warming,  a  few  drops  of 
ether,  which  will  expel  the  air  and  prevent  a  possible  explosion. 
The  phosphine  prepared  in  this  manner  contains  some  hydrogen 
and  some  of  the  liquid  hydrogen  phosphide,  P2H4,  which  corres- 
ponds in  composition  to  hydrazine,  N2H4.  This  liquid  is  volatile 
and  takes  fire  spontaneously  on  exposure  to  the  air.  For  this 
reason,  although  the  kindling  temperature  of  phosphine  is  about 
150°,  the  phosphine  prepared  as  described  takes  fire  at  once  as 
it  comes  to  the  air.  Bubbles  of  the  gas  explode  as  they  come  to 
the  surface  of  the  water,  forming  a  cloud  of  phosphoric  acid, 
H3PO4,  which  gives  beautiful  vortex  rings  in  still  air  (Fig.  78). 
Phosphine  may  be  condensed  to  a  colorless  liquid,  which  boils 
at  —  86.2°  and  solidifies  at  lower  temperatures  to  crystals  which 
melt  at  -  133°. 

Phosphonium  Salts.  Phosphine  combines  with  acids  to  form 
phosphonium  salts,  as  ammonia  forms  ammonium  salts.  The 
most  stable  and  best  known  of  these  salts  is  phosphonium  iodide, 
PH4I,  which  may  be  prepared  by  the  direct  union  of  phosphine, 
PH3,  and  hydriodic  acid,  HI.  It  forms  white  crystals,  which 


244 


A  TEXTBOOK  OF  CHEMISTRY 


sublime  at  80°.  The  salt  is  hydrolyzed  by  water  and  phosphine 
escapes  from  the  solution  : 

PH4I  +  HOH  =  PH4OH  +  HI 
PH4OH  =  PH3  +  HOH 

Evidently  the  phosphonium  group,  PH4,  is  very  unstable,  even 
in  the  presence  of  hydrogen  ions.  It  is  sometimes  stated  that 
phosphine  is  a  much  weaker  base  than  ammonia.  Correctly 

speaking  neither  is 
a  base,  and  the  true 
base,  phosphonium 
hydroxide,  PH4OH, 
is  extremely  un- 
stable, if  it  exists 
at  all.  We  shall 
find  that  arsine, 
AsH3,  and  stibine, 
SbH3,  do  not  com- 
bine with  acids. 
In  the  series  NH3, 
PH3,  AsH3,  SbH3, 
not  only  does  the 
tendency  to  com- 
bine with  a  fourth 
hydrogen  atom  be* 

come  less  and  less,  but  the  compounds  themselves  are  less  and 
less  stable,  stibine,  SbH3,  decomposing  at  ordinary  temperatures, 
especially  in  the  presence  of  metallic  antimony. 

Phosphorus  Trichloride,  PC13,  and  Phosphorus  Pentachloride, 
PCls,  are  easily  prepared  by  the  direct  union  of  chlorine  and 
phosphorus.  The  trichloride  is  a  liquid  which  boils  at  76°. 
The  pentachloride  is  a  white  solid  which  melts  in  a  sealed  tube 
at  148°.  Its  vapor  pressure,  however,  is  760  mm.  at  140°.  In 
other  words  its  melting  point  is  higher  than  its  boiling  point  and 
it  sublimes  without  melting  when  heated  under  atmospheric 
pressure. 


Fig.  78 


CHLORIDES  OF  PHOSPHORUS  245 

The  molecular  weight  of  phosphorus  pentachloride,  PC15, 
is  208.5  ;  but  a  gram  molecular  volume  of  the  gas  at  182°  weighs 
147  grams,  while  at  300°  it  weighs  only  105.7  grams,  only  a  little 
more  than  one  half  the  weight  of  a  gram  molecule.  This  indi- 
cates that  the  pentachloride  dissociates  into  phosphorus  trichlo- 
ride and  chlorine,  the  dissociation  being  nearly  complete  at  300°. 
This  gives  twice  as  many  molecules  as  there  are  in  the  original 
pentachloride  and  one  gram  molecule  of  the  pentachloride  gives 
two  gram  molecular  volumes  of  gas  : 

PC15  ^±  PC13  +  C12 

Hydrolysis  of  the  Chlorides  of  Phosphorus.  The  chlorides 
of  phosphorus  are  decomposed,  or  hydrolyzed  by  water  in  the 
same  manner  as  most  other  chlorides  of  nonmetallic  elements  : 

PC13  +  3H.OH  =   H3P03  +  3HC1 

Phosphorous 
Acid 

PC16  +  4H.OH  =  H3PO4  +  5HC1 

Phosphoric 
Acid 

Phosphorus  Oxychloride,  POC13.  When  phosphorus  penta- 
chloride is  treated  with  a  small  amount  of  water  or  with  almost 
any  compound  containing  the  hydroxyl  group,  OH,  it  is  changed 
to  phosphorus  oxychloride,  POC13,  while  the  two  chlorine  atoms 
which  are  lost  combine  with  the  two  atoms  which  were  united 
to  the  oxygen  : 

cl 


x    / 

.P^CI  +    HOH     = 
c/  xci  xci  . 


an  Ethyl  Alcohol 

Ethyl  Chloride 

a       a  a 

;p^a  +    c2H4o2     = 

Cl'     XC1      (orC2H3O.OH) 

Acetic  Acid 


246  A  TEXTBOOK  OF  CHEMISTRY 

Phosphorus  oxychloride  may  also  be  prepared  by  oxidizing 
phosphorus  trichloride  with  potassium  chlorate.  It  is  a  color- 
less liquid  which  boils  at  107.2°.  It  has  a  very  unpleasant  odor 
and  the  vapor  attacks  the  eyes  strongly.  It  is,  of  course,  hy- 
drolyzed  by  water  to  phosphoric  and  hydrochloric  acids.  It  is 
the  chloride  of  phosphoric  acid,  H3PC>4,  in  the  same  sense  that 
sulfuryl  chloride,  SC^Ck,  is  the  chloride  of  sulfuric  acid,  H2S04 
(p.  189). 

Oxides  of  Phosphorus.  When  phosphorus  is  burned  with  an 
insufficient  supply  of  air,  a  mixture  of  two  oxides,  phosphorus 
"  trioxide,"  P^e,  and  phosphorus  "  pentoxide,"  P4Oio,  is  formed. 
The  names  were  given  long  before  determinations  of  the  density 
of  the  vapors  of  these  compounds  showed  that  they  have  the 
formulas  given.  The  names  refer,  of  course,  to  the  simple  formu- 
las P2O3  and  P2O6.  The  trioxide,  PA,  is  a  solid  which  melts  at 
22.5°  and  boils  at  173.1°.  The  pentoxide  is  a  solid  which  may  be 
sublimed  at  a  high  temperature  but  which  gives  off  almost  no 
vapor  at  ordinary  temperatures.  It  has  a  very  strong  affinity 
for  water  and  is  the  most  perfect  drying  agent  for  gases  which  we 
have.  If  10,000  liters  of  air  are  passed  through  a  comparatively 
small  tube  filled  with  the  pentoxide,  no  moisture  which  can  be 
determined  remains  in  the  gas,  while  the  vapor  of  the  pentoxide 
which  is  carried  away  by  the  gas  weighs  only  one  milligram  (Mor- 
ley,  Am.  J.  Sci.  34,  199  (1887) ;  J.  Am.  Chem.  Soc.26,  1171  (1904). 

Another  oxide  of  phosphorus,  ?2O4,  called  phosphorus  tetrox- 
ide,  is  known,  but  is  of  little  interest  except  as  corresponding  to 
nitrogen  tetroxide,  N2O4. 

Acids  of  Phosphorus.  While  the  acids  of  nitrogen,  nitrou 
acid,  HNO2,  and  nitric  acic^  HNO3,  are  formed  by  the  addition 
of  one  molecule  of  water  to  the  anhydrides,  N2O3  and  N2O5, 
the  normal  acids  of  phosphorus  corresponding  to  these  are 
formed  by  the  addition  of  three  molecules  of  water  to  the  cor- 
responding anhydrides  (using  the  simpler  formulas).  Two  other 
acids,  metaphosphoric  acid,  HPO3,  and  pyrophosphoric  acid, 
H4P2O7,  are  also  derived  from  phosphorus  pentoxide  (phosphoric 
anhydride),  P2O5.  As  the  addition  of  water  is  considered  as 


ACIDS  OF  PHOSPHORUS  247 

neither  an  oxidation  nor  a  reduction,  the  three  acids  derived 
from  the  pentoxide  are  all  called  "  phosphoric  "  acids  and  are 
distinguished  by  prefixes.  These  relations  will  be  clearer  from 
the  following  table : 

(P2O1.3H2O)=  2  H3PO2  Hypophosphorous  acid 

(P2O3.3  H2O)  =  2  H3PO3  Phosphorous  acid 

(P2O6.3  H2O)  =  2  H3PO4  Orthophosphoric  acid  ' 

(P2O5.2  H2O)  =  H4P2O7  Pyrophosphoric  acid 

(P2O5.H2O)     =  2  HPO3  Metaphosphoric  acid 

Basicity  of  the  Acids  of  Phosphorus.  The  formulas  of  hypo- 
phosphorous  acid,  H3PO2,  phosphorous  acid,  H3PO3,  and  ortho- 
phosphoric  acid,  H3PO4,  might  lead  us  to  expect  each  of  these 
acids  to  be  tribasic.  It  is  found,  however,  that  only  one  atom  of 
hydrogen  in  hypophosphorous  acid  can  be  replaced  by  metals 
and  only  two  of  the  hydrogen  atoms  in  phosphorous  acid.  In 
1  other  words  hypophosphorous  acid  is  monobasic,  phosphorous 
acid  dibasic,  and  orthophosphoric  acid  tribasic.  The  normal 
sodium  salts  are : 

I  Sodium  hypophosphite,     NaH2PO2 

Sodium  phosphite,  Na2HPO3 

Sodium  orthophosphate,   Na3PO4 

These  and  other  facts,  which  cannot  be  given  here,  make  it 
probable  that  the  structure  of  these  acids  is  correctly  represented 
by  the  following  formulas  : 

Hypophosphorous  acid 

H-       X)-H 

H-0         6 

Phosphorous  acid  /P  / 

W 

H-0          O 

i  Orthophosphoric  acid  /P\ 

K—O' 


1  This  oxide  is  given  in  many  of  the  books,  but  its  existence  is 
extremely  doubtful. 


248  A  TEXTBOOK  OF  CHEMISTRY 

According  to  .these  formulas  only  the  hydrogen  atoms  which 
are  united  to  oxygen  are  acid  in  character.  Also  the  oxidation 
of  the  lower  acids  consists  in  the  introduction  of  an  oxygen  atom 
between  a  hydrogen  atom  and  the  phosphorus.  According 
to  the  electron  theory  the  valence  of  the  phosphorus  is  negative 
toward  the  hydrogen  atoms  and  positive  towards  the  oxygen. 
Oxidation,  in  such  a  case,  consists  in  the  change  of  a  negative 
valence  to  a  positive  one. 

Hypophosphorous  Acid,  H3PO2.  The  sodium  salt  of  hypo- 
phosphorous  acid,  NaH2PO2,  is  formed  when  phosphorus  is 
warmed  with  a  solution  of  sodium  hydroxide,  phosphine,  PH3, 
being  evolved  at  the  same  time.  The  acid  is  monobasic.  It  is 
a  powerful  reducing  agent.  Some  hypophosphites  are  used 
in  medicine. 

Phosphorous  Acid,  H3PO3,  is  formed  with  phosphoric  and 
hypophosphoric  acids,  when  ordinary  phosphorus  is  allowed  to 
oxidize  slowly  in  moist  air,  but  it  is  extremely  difficult  to  separate 
the  mixture  into  its  components.  The  pure  acid  may  be  pre- 
pared by  the  hydrolysis  of  phosphorus  trichloride.  It  is  a  bi- 
basic  acid,  the  two  sodium  salts  being  monosodium  phosphite, 
NaH2PO3,  and  disodium  phosphite,  Na2HP03.  Phosphorous 
acid  is  also  a  powerful  reducing  agent. 

Orthophosphoric  Acid,  H3PC>4,  is  formed  when  phosphoric 
anhydride,  P4Oio,  is  dissolved  in  hot  water.  It  is  also  formed 
when  solutions  of  pyrophosphoric  acid,  H4P2O7,  or  metaphos- 
phoric  acid,  HPO3,  are  boiled,  especially  if  some  strong  acid,  as 
nitric  acid  or  hydrochloric  acid,  is  present  to  catalyze  the  reac- 
tion, which  is  to  be  considered  as  a  hydrolysis : 


OH  OH 


+     HOH 


Pyrophosphoric  Acid  Orthophosphoric  Acid  (2  mols) 


PHOSPHORIC  ACIDS  249 

HOH    =    O= 


XOH 

Metaphosphoric  Acid 

Pure  orthophosphoric  acid  forms  clear,  rhombic  crystals, 
which  melt  at  about  40°.  These  crystals  dissolve  in  a  small 
amount  of  water,  forming  a  heavy,  sirupy  liquid  somewhat 
resembling  concentrated  sulfuric  acid  in  appearance. 

An  impure  solution  of  phosphoric  acid  was  formerly  prepared 
on  a  large  scale,  as  a  step  in  the  manufacture  of  phosphorus,  by 
treating  bone  ash  with  dilute  sulfuric  acid  and  filtering  the  solu- 
tion from  the  calcium  sulfate,  which  is  only  slightly  soluble  in 
water  : 

Ca3(PO4)2  +  3  H2SO4  =  2  H3PO4  +  3  CaSO4 

Tricalcium  Calcium 

Phosphate  Sulfate 

Orthophosphoric  acid  forms  three  classes  of  salts,  in  which 
one,  two  or  three  atoms  of  hydrogen  are  replaced  in  each  mole- 
cule of  the  acid.  These  are  called  primary,  secondary  and  ter- 
tiary, or,  more  often,  mono-,  di-  and  tri-metallic  salts.  The 
following  are  the  names  of  the  sodium  and  calcium  salts  : 

Monosodium  phosphate,  NaH2PO4  (primary) 
Monocalcium  phosphate,  Ca(H2PO4)2  (primary) 
Disodium  phosphate,         Na2HPO4  (secondary) 
Dicalcium  phosphate,        CaHPO4  (secondary) 
Trisodium  phosphate,        Na3PO4  (tertiary) 
Tricalcium  phosphate,       Cas(PO4)2  (tertiary) 

Orthophosphoric  acid  is  much  the  most  important  of  the  acids 
of  phosphorus,  being  the  acid  into  which  all  of  the  others  tend 
to  pass  either  by  oxidation  or  hydrolysis.  Apart  from  its  oc- 
currence in  organic  compounds,  phosphorus  is  found  almost 
exclusively  in  the  form  of  orthophosphates,  and  these  phosphates 
are  an  indispensable  constituent  of  arable  soils. 


250  A  TEXTBOOK  OF  CHEMISTRY 

lonization  of  Orthophosphoric  Acid.  Orthophosphoric  acid 
is  an  acid  of  only  moderate  strength.  A  TV  formular  (i.e.  con- 
taining one  gram  molecule  in  10  liters  of  water)  solution  contains, 
of  course,  three  times  as  many  replaceable  hydrogen  atoms  as 
a  tenth  normal  solution  of  hydrochloric  acid,  but  it  contains  only 
one  third  as  many  hydrogen  ions.  This  means  that  in  the  ioniza- 
tion  reaction  : 


the  equilibrium  is  comparatively  far  to  the  left,  even  in  quite 
dilute  solutions.  Even  in  very  dilute  solutions  the  second  and 
third  hydrogen  atoms  ionize  to  only  a  very  slight  extent.  This 
may  be  either  because  the  three  hydrogen  atoms  in  Orthophos- 
phoric acid  are  different  or  because  after  the  removal  of  the  one 
hydrogen  atom  the  negative  ion,  H2PO4~,  holds  the  remaining 
hydrogen  atoms  too  strongly  for  them  to  separate  easily  as  ions. 
The  second  explanation  seems  more  probable.  If  the  first  hy- 
drogen atom  is  completely  neutralized  by  the  addition  of  a  base  : 

H+  +  H2P04-  +  Na+  +  OH-  ^±  Na+  +  H2PO4-  +  H2O 

the  dihydrogen  phosphate  ion,  H2PO4~,  will  ionize  to  a  slight 
extent : 

H2P04~  ^. 


but  the  solution  is  only  faintly  acid,  and  if  more  sodium  hydroxide 
is  added  to  'such  a  solution,  the  accumulation  of  the  monohy- 
drogen  phosphate  ions,  HPO4=,  shifts  the  equilibrium  to  the  left. 
This  is  because  sodium  salts  of  weak  acids  are  always  much 
more  completely  ionized  than  the  corresponding  acids.  This 
shifting  of  the  equilibrium  prevents  much  formation  of  new 
hydrogen  ions,  as  those  which  are  present  are  removed  by  com- 
bination with  the  hydroxide  ions  of  the  sodium  hydroxide. 
Before  all  of  the  second  hydrogen  atoms  of  the  phosphoric  acid 
have  been  neutralized,  the  tendency  of  the  monohydrogen  phos- 
phate ions,  HPO4=,  to  combine  with  hydrogen  ions  will  become  so 
strong  that  even  the  hydrogen  ions  of  water  will  combine  with 
them,  leaving  an  excess  of  hydroxide  ions  in  the  solution.  Such 
a  solution  must,  of  course,  react  alkaline.  From  this  conduct 


IONIZATION  OF  PHOSPHORIC  ACID  251 


of  phosphoric  acid  it  is  evident  that  if  we  attempt  to  titrate  a 
solution  of  phosphoric  acid  by  adding  sodium  hydroxide,  instead 
of  the  sharp  change  which  occurs  in  titrating  hydrochloric  or 
sulfuric  acid,  there  will  be  a  gradual  change  from  a  solution  con- 
taining a  slight  excess  of  hydrogen  ions,  H+,  to  one  containing 
a  slight  excess  of  hydroxide  ions,  OH~.  Two  things  result  from 
these  properties  of  solutions  containing  salts  of  phosphoric  acid  : 
first,  unless  a  very  sensitive  indicator  is  chosen,  that  is,  one  in 
which  the  change  in  color  is  produced  by  a  very  slight  change 
in  the  ratio  between  the  hydrogen  and  hydroxide  ions  present, 
the  end  point  of  the  titration  will  be  indefinite ;  and,  second, 
since  most  indicators  change  color,  not  when  the  number  of 
hydrogen,  H+,  and  hydroxide,  OH~,  are  equal,  but  when 
there  is  an  excess  of  one  or  the  other,  and  this  excess 
differs  for  different  indicators,  the  end  point  in  titrating 
phosphoric  acid  will  depend  on  the  indicator  chosen.  (See 
p.  387.)  Thus,  if  methyl  orange  or  cochineal  is  used,  the 
end  point  in  fairly  dilute  solutions  will  be  found  when 
the  solution  corresponds  very  nearly  to  the  composition 
NaH2PO4.  With  phenolphthalein,  on  the  other  hand,  the  end 
will  correspond  nearly  to  the  composition  Na2HPO4.  With 
litmus  the  end  lies  between  the  two.  If  alizarine  green  is  used, 
the  change  in  color  occurs  when  the  composition  of  the  solution 
is  very  nearly  represented  by  the  formula  Na3PO4. 

It  is  well,  also  to  consider  the  conduct  of  disodium  phosphate 
from  a  somewhat  different  point  of  view,  which,  however,  follows 
logically  from  what  has  been  said.  If  the  salt,  which  crystallizes 
with  the  composition  Na2HPO4.12  H2O,  is  dissolved  in  water,  we 
should  expect  the  formation  of  the  ions,-  Na+  +  Na+  +  HPO4=. 
But,  as  has  been  stated,  in  the  presence  of  many  of  the  mono- 
hydrogen  phosphate  ions,  HPO4=,  these  have  a  tendency  to  form, 
with  the  hydrogen  ions  of  the  water,  dihydrogen  phosphate 
ions,  H2PO4~,  because  the  latter  ionize  only  to  a  slight  extent. 
This  results  in  the  presence  of  an  excess  of  hydroxide  ions  in  the 
solution,  which  will  react  alkaline  toward  indicators  that  are 
sensitive  to  a  slight  excess  of  hydroxide  ions. 


252  A  TEXTBOOK  OF  CHEMISTRY 


Na+  +  Na+  +  HPO4=  +  H+ 

Water  in 
Ionic  Form 

^±  Na+  +  Na+  +  H2PO4 

This  sort  of  hydrolysis  occurs  with  all  salts  of  strong  bases, 
as  sodium  hydroxide  or  potassium  hydroxide,  with  weak  acids 
or  with  acids  whose  second  or  third  hydrogen  atoms  undergo 
slight  ionization. 

Trisodium  phosphate,  Na3PO4,  will,  of  course,  be  much  more 
completely  hydrolyzed  : 

Na+  +  Na+  +  Na+  +  PO4^  +  H+  +  OH~ 

^t  3  Na+  +  HP04=  +  OH- 

The  only  tertiary  salts  of  orthophosphoric  acid  which  are  sol- 
uble in  water  are  those  of  the  alkali  metals,  sodium,  potassium, 
etc.  All  other  tertiary  or  normal  phosphates  are  insoluble. 
Many  of  the  primary  and  secondary  phosphates  are  either 
insoluble  or  are  decomposed  by  water  into  phosphates  which 
approach  the  tertiary  phosphates  in  composition,  and  either 
phosphoric  acid  or  more  acid  phosphates,  which  dissolve  in  an 
excess  of  the  acid. 

Decomposition  of  Primary  and  Secondary  Salts  of  Ortho- 
phosphoric  Acid.  Salts  of  orthophosphoric  acid  which  contain 
hydrogen  decompose  on  heating,  losing  all  of  their  hydrogen  as 
water  and  leaving  salts  of  metaphosphoric  or  pyrophosphoric 
acid.  As  ammonium  salts  dissociate  on  heating,  these  give  the 
same  products  as  if  they  contained  hydrogen  in  place  of  ammo- 
nium, NH4.  Monosodium  phosphate,  NaH2PO4,  gives  sodium 
metaphosphate,  NaPO3;  and  sodium  ammonium  phosphate, 
NaNH4HPO4,  gives  the  same  compound.  Disodium  phosphate, 
Na2HPO4,  gives  sodium  pyrophosphate,  Na4P2O7;  and  am- 
monium magnesium  phosphate,  NH4MgPO4,  gives  magnesium 
pyrophosphate,  Mg2P2O7.  Magnesium  diammonium  phosphate, 
Mg(NH4)4(PO4)2,  or  Mg  [(NH4)2PO4]2,  gives  magnesium  meta- 
phosphate, Mg(PO3)2. 


PHOSPHORIC  ACIDS  253 

Pyrophosphoric  Acid,  H4P2O7.  If  orthophosphoric  acid  is 
heated  carefully  at  250°  it  loses  water  and  is  changed  to  pyro- 
phosphoric  acid : 

2H3PO4-H20  =  H4P2O7 

The  acid  may  be  dissolved  in  cold  water,  giving  a  solution  which 
differs  in  its  properties  from  those  of  a  solution  of  orthophos- 
phoric acid.  Especially,  after  neutralization  it  gives  with  silver 
nitrate,  AgNOs,  a  white  precipitate  of  silver  pyrophosphate, 
Ag4P2O7,  while  orthophosphoric  acid,  or  orthophosphates, 
will  give  a  yellow  precipitate  of  trisilver  phosphate,  AgsPC^. 
Sodium  pyrophosphate  is  easily  prepared  by  heating  disodium 
phosphate,  Na2HPO4. 

Metaphosphoric  Acid,  HPO3,  is  formed  when  phosphoric 
anhydride,  P4Oio,  is  dissolved  in  cold  water  or  when  either  ortho- 
phosphoric  acid  or  pyrophosphoric  acid  is  heated  to  a  high  tem- 
perature. It  differs  from  the  other  two  phosphoric  acids  in  that 
its  neutralized  solution  gives  with  silver  nitrate  a  white  pre- 
cipitate of  silver  metaphosphate,  AgPO3,  instead  of  the  yellow 
precipitate  of  trisilver  phosphate,  Ag3PO4,  given  by  ortho- 
phosphoric  acid  and  the  white  precipitate  of  silver  pyrophos- 
phate, Ag4P2O7,  given  by  pyrophosphoric  acid.  Metaphos- 
phoric acid  also  precipitates  a  solution  of  albumin,  as  of  the 
white  of  an  egg,  which  has  been  acidified  with  acetic  acid,  while 
ortho-  and  pyrophosphoric  acids  or  their  salts  do  not  do  this. 

Sodium  metaphosphate  may  be  prepared  by  heating  either 
monosodium  phosphate,  NaH2PO4,  or  sodium  ammonium  phos- 
phate, NaNH4HPO4.  This  last  salt  is  called  microcosmic  salt 
and  is  used  in  blowpipe  analysis.  When  this  salt  is  heated  in 
a  loop  of  platinum  wire,  it  melts  to  a  clear  bead  of  sodium  meta- 
phosphate, NaPO3,  which  will  dissolve  the  oxides  of  many  of 
the  metals,  forming  double  salts  of  orthophosphoric  acid : 

NaPO3  +  CuO  =     NaCuPO4 

Sodium  Copper 
Orthophosphate 


254  A  TEXTBOOK  OF  CHEMISTRY 

The  metaphosphate  may  be  considered  here  as,  in  a  certain 
sense,  an  acid  anhydride  which  with  oxides  forms  normal  salts  of 
orthosphoric  acid.  The  copper  sodium  phosphate  is  blue,  and 
the  colors  given  to  the  microcosmic  bead  by  different  metallic 
oxides  serve  as  a  means  for  their  identification. 

*  Metaphosphoric  acid  may  be  vaporized  at  a  high  tempera- 
ture and  the  vapor  has  the  formula  (HPOs^.     A  study  of  the 
salts  and  of  the  properties  of  solutions  of  the  acid  prepared  in 
different  ways  has  shown  that  several  polymeric  forms  of  the 
acid  exist,  that  is,  forms  having  the  same  composition  but  dif- 
ferent molecular  weights.     The  salts  of  these  various  forms  are 
called  dimetaphosphates,  M2P2Oe,  trimetaphosphates,  M3P3O9, 
tetrametaphosphates,  M4P4Oi2,  etc.     In  these  formulas  "  M  " 
is  used  to  represent  any  univalent  metal. 

*  Hypophosphoric  Acid,  H2PO3,  is  one  of  the  products  formed 
by  the  slow  oxidation  of  phosphorus  in  moist  air.     From  the 
solution  obtained  in  this  manner  the  rather  difficultly  soluble 
acid  sodium  salt,  NaHPO3,  is  precipitated  by  a  concentrated 
solution  of  sodium  acetate,  NaC2H3O2. 

From  its  formula  we  should  expect  that  phosphorus  tetroxide, 
P2O4,  would  be  the  anhydride  of  hypophosphoric  acid,  but, 
curiously  enough,  when  phosphorus  textroxide  is  dissolved  in 
water,  a  mixture  of  phosphorous  and  orthophosphoric  acids  is 
formed : 

P2O4  +  3  H2O  =  H3PO3  +  H3P04 

The  true  anhydride  of  hypophosphoric  acid  (PO2?)  has  not 
been  prepared. 

*  Sulfides  of  Phosphorus.     Four  sulfides  of  phosphorus  have 
been  prepared,  tetraphosphorus  trisulfide,  P4S3,  tetraphosphorus 
heptasulfide,  P4S7,  triphosphorus  hexasulfide,  P3S6,  and  diphos- 
phorus  pentasulfide,  P2S5.     The  last  is  usually  called  phosphorus 
pentasulfide.     It  melts  at  274°-276°  and  boils  at  530°.     It  has 
been  used  in  chemical  laboratories  frequently  to  obtain  a  nearly 
constant,  high  temperature. 

Tetraphosphorus  trisulfide,  P4S3,  melts  at  165°-166°  and  boils 


PHOSPHORUS  255 

at  225°-235°  under  a  pressure  of  10  mm.  As  it  takes  fire  with 
slight  friction  and  as  its  vapors  are  either  nonpoisonous,  or,  in 
any  case,  far  less  poisonous  than  those  of  ordinary  phosphorus, 
it  is  likely  to  replace  the  latter  entirely  for  the  manufacture  of 
matches. 

EXERCISES 

1.  Write  the  equation  for  the  reaction  between  phosphorus  and  a 
solution  of  sodium  hydroxide,  giving  hydrogen  and  sodium  hypophos- 
phite. 

2.  Write  the  equation  for  the  reaction  between  phosphorus  and 
sodium  hydroxide  in  solution,  giving  phosphine  and  sodium  hypophos- 
phite. 

3.  Write  the  equation  for  the  reaction  giving  liquid  hydrogen  phos- 
phide, P2Hj,  and  sodium  hypophosphite. 

4.  What  is  the  distinction  between  a  substance  which  sublimes  and 
one  which  boils  ?     Under  what  conditions  does  water  sublime  ? 

5.  What  percent  of  phosphorus  pentachloride  is  dissociated  when  its 
gram  molecular  volume  weighs  156.4  grams  ?     What  per  cent  when  it 
weighs  130  grams  ? 

6.  If  on  heating  phosphorous  and  hypophosphorous  acids  the  prod- 
ucts formed  are  phosphine,  metaphosphoric  acid  and  water,  what  are 
the  equations  representing  the  decomposition  of  these  acids?     Are 
these  decompositions  consistent  with  the  structural  formulas  which 
have  been  proposed  for  these  acids  ? 

7.  If  the  structure  of  phosphorous  acid  were  correctly  represented  by 

X)— H 

the  formula  P\-O — H,  how  ought  it  to  decompose  on  heating  ? 
X0-H     • 

8.  Metaphosphoric  acid  volatilizes  at  a  very  much  higher  tempera- 
ture than  sulfuric  acid.     What  will  be  the  effect  of  heating  a  mixture 
of  sodium  sulf  ate  and  metaphosphoric  acid  ? 

9.  How  much  iodine  and  water  will  be  required  to  convert  10  grams 
of  phosphorus  into  orthophosphoric  acid  if  the  reaction  is  quantitative  ? 

10.  How  many  liters  of  air  will  be  required  to  burn  10  grams  of 
phosphorus  to  the  pentoxide  ? 


CHAPTER  XV 
ARSENIC,   ANTIMONY   AND   BISMUTH 

IT  has  been  pointed  out  that  with  increasing  atomic  weights 
the  elements  of  the  nonmetallic  groups  of  the  Periodic  System 
become  more  metallic  in  character.  This  is  especially  evident 
in  the  fifth  group.  Arsenic  is  metallic  in  its  appearance,  opaque 
and  like  steel  on  its  surface  when  not  tarnished.  It  is,  however, 
brittle,  and  its  chloride,  AsCl3,  is  hydrolyzed  by  water,  resembling 
the  chlorides  of  the  nonmetals  rather  than  those  of  the  metals. 
Arsenic  forms  no  salts  with  sulfuric,  nitric  or  other  acids.  Anti- 
mony and  bismuth  are  still  more  metallic  in  their  appearance 
and  bismuth  is  malleable  to  a  slight  extent.  Their  chlorides  are 
hydrolyzed  by  water,  at  first,  only  to  the  oxychlorides,  SbOCl 
and  BiOCl.  Both  of  them  form  normal  nitrates,  Sb(NO3)3,  and 
Bi(NO3)3,  and  sulfates,  Sb2(SO4)3,  Bi(SO4)3,  though  these  are 
hydrolyzed  to  basic  salts  or  even  to  the  hydroxides  or  oxides  by 
water. 

Arsenic,  As,  74.96.  Occurrence.  Arsenic  is  found  in  the  free 
state  in  nature,  but  occurs  chiefly  combined  with  sulfur,  either 
alone,  as  in  the  disulfide,  realgar,  As2S2,  or  tiie  trisulfide,  orpiment, 
As2S3,  or,  much  more  frequently,  with  the  sulfides  of  other 
metals,  the  most  common  compound  of  this  kind  being  arseno- 
pyrite,  or  mispickel,  FeAsS.  Iron  pyrites  and  copper  pyrites 
almost  invariably  contain  arsenic,  often  in  considerable  quan- 
tities. From  the  former  the  arsenic  finds  its  way  into  commer- 
cial sulfuric  acid  and  from  that  into  a  great  variety  of  chemical 
products.  From  the  copper  pyrites  the  arsenic  escapes  along 
with  the  sulfur  dioxide  in  the  process  of  roasting,  no  less  than 
twenty-five  tons  a  day  of  arsenic  trioxide  escaping  from  a  single 
smelting  furnace  in  Montana  (J.  Am.  Chem.  Soc.  29,  993  (1907). 

256 


ARSENIC 


257 


Preparation  and  Properties  of  Arsenic.  Metallic  arsenic  is 
usually  prepared  by  heating  arsenopyrite,  FeAsS,  the  arsenic 
subliming  and  leaving  ferrous  sulfide,  FeS,  behind.  Prepared  in 
this  manner  it  forms  a  dark  gray,  brittle  mass.  Fragments 
heated  in  a  closed  tube  or  before  the  blowpipe  on  charcoal,  so 
that  the  tarnished  surface  is  removed,  appear  like  steel.  When 
deposited  on  a  glass  or  porcelain  surface  (Marsh's  test),  arsenic 
is  brown  or  black  according  to  the  thickness  of  the  deposit, 
usually  showing  brown  at  the  edges  where  the  deposit  is  thin, 
while  antimony  is  a  more  sooty  black.  Here,  again,  we  have  an 
increase  in  metallic  properties  with  increasing  atomic  weight, 
opacity  being  one  of  the  most  marked  properties  of  metals. 

The  specific  gravity  of  gray  arsenic  is  5.73.  The  formula  of 
its  vapor  at  560°-670°  is  As4,  at  1700°,  As2.  The  melting  point 
of  arsenic  is  higher  than  its  boiling  point,  hence  it  sublimes  with- 
out melting  when  heated  on  charcoal  or  in  a  tube  closed  at  one 
end,  a  property  which  distinguishes  it  easily  from  antimony. 

Metallic  arsenic  is  sometimes  used  for  poisonous  fly  papers. 
Three  one-hundredths  of  a  per  cent  of  arsenic  lowers  the  con- 
ductivity of  copper  14  per  cent  and  injures  it  seriously,  especially 
for  electrical  use. 

Arsine,  AsH3.  Marsh's  Test.  When  almost  any  soluble 
compound  of  arsenic  is  added  to  a  flask  in  which  hydrogen  is 
being  generated  from 
zinc  and  sulfuric  or  hy- 
drochloric acid,  the 
arsenic  is  reduced  to 
arsine.  If  the  hydrogen 
containing  arsine  is  con- 
veyed through  a  hard 
glass  tube,  narrowed  at 
one  point  (Fig.  79),  and 


Fig.  79 


the  tube  is  heated  just  back  of  the  constriction  with  a  Bunsen 
flame,  the  arsine  is  decomposed  and  metallic  arsenic  is  deposited 
as  a  brown  or  black  mirror  on  the  glass.  As  small  a  quantity 
of  arsenic  as  ytfW  °f  a  milligram  can  be  seen  in  this  form, 


258  A  TEXTBOOK  OF  CHEMISTRY 

and  the  process  has  been  long  used,  under  the  name  of  Marsh's 
test,  for  the  detection  and  estimation  of  small  quantities  of 
arsenic,  especially  in  cases  of  criminal  poisoning,  or  for  the 
examination  of  wall  papers  or  articles  of  food.  One  of  the 
first  requisites  in  making  the  test  is,  of  course,  that  the  zinc, 
sulfuric  acid  and  other  materials  used  should  be  entirely  free 
from  arsenic.  Commercial  zinc  and  commercial  sulfuric  acid 
almost  invariably  contain  the  element.  It  is  necessary,  also, 
to  distinguish  the  mirror  from  that  of  antimony,  which  may  be 
obtained  in  the  same  manner.  When  the  amount  of  arsenic 
is  considerable,  it  imparts  to  the  burning  hydrogen  flame  a  pale 
blue  color,  and  arsenic  is  deposited  on  a  piece  of  porcelain  held 
in  the  flame,  very  much  as  soot  is  deposited  from  a  candle 
flame. 

Arsine  may  be  condensed  to  a  liquid,  which  boils  at  —  55°. 
It  is  very  poisonous.  Some  years  ago  a  chemist  in  Chile  was 
fatally  poisoned  while  working  with  it. 

Arsine  does  not  show  any  tendency  to  combine  with  acids,  as 
ammonia  and  phosphine  do. 

Arsenic  "  Trioxide,"  As4O6.  When  arsenic  is  heated  in  the 
air,  it  burns  to  arsenic  "  trioxide,"  frequently  called  white 
arsenic.  The  simpler  formula,  As2Os,  is  commonly  used  for  the 
compound,  but  the  density  of  its  vapor  corresponds  to  the 
formula  As^e.  It  crystallizes  in  octahedra  which  are  highly 
characteristic,  and  the  microscopic  identification  of  the  crystals 
is  one  of  the'  most  important  means  of  demonstrating  the  pres- 
ence of  arsenic. 

Arsenic  trioxide  is  one  of  the  most  common  compounds  of  the 
element  and  is  frequently  used  as  a  ratsbane  and  has  often 
been  used  for  criminal  poisoning.  The  fatal  dose  for  an  adult  is 
from  0.06  to  0.18  gram  (one  to  three  grains),  but  it  seems  possible 
to  accustom  the  organism  to  its  use,  and  the  so-called  arsenic 
eaters  may  sometimes  take  four  times  that  amount  without 
apparent  injury.  The  best  antidote  is  freshly  precipitated 
ferric  hydroxide,  Fe(OH)3,  or  a  colloidal  solution  of  ferric 
hydroxide. 


ARSENIC  259 

Crystallized  arsenic  trioxide  dissolves  in  50  parts  of  water  at 
25°.  The  amorphous  form  is  somewhat  more  soluble.  The 
solution  reacts  faintly  acid,  and  forms  salts  with  bases,  but  on 
evaporation  it  deposits  the  trioxide. 

Arsenious  Acid.  As  has  just  been  stated,  arsenious  acid 
resembles  sulfurous  and  nitrous  acids  in  that  it  exists  only  in 
solution  and  decomposes  easily  into  its  anhydride  and  water. 
Salts  of  acids  derived  from  this  anhydride  are  known,  however. 
Among  these  may  be  mentioned  silver  orthoarsenite,  Ag3AsO3, 
and  monopotassium  diarsenite,  KHAS2O4.  The  last  seems  to 
be  derived  from  a  diarsenious  acid,  H2As2O4,  which  would  cor- 
respond to  a  doubled  nitrous  acid,  (HNO2)2-  Paris  green  is  a 
double  salt  of  copper  with  acetic  and  arsenious  acids, 
Cu(C2H302)2.Cu3(As03)2. 

Arsenic  Pentoxide,  As2O5,  and  Arsenic  Acid,  H3AsO4.  When 
arsenic  trioxide  is  warmed  with  nitric  acid,  it  is  oxidized  to 
arsenic  acid : 

2  HNO3  +  As2O3  +  2  H20  =  2  H3AsO4  +  NO  +  NO2 

From  a  concentrated  solution  the  acid  crystallizes  with  one 
molecule  of  water,  H3AsO4.H2O.  At  140°-180°  this  hydrate 
loses  water  and  gives  pyroarsenic  acid,  H4As2C>7,  and  at  200° 
the  latter  loses  more  water  and  gives  metarsenic  acid,  HAsO3. 
At  a  higher  temperature  metarsenic  acid  loses  more  water  and 
arsenic  pentoxide,  As2Os,  remains.  This  cannot  be  volatilized 
without  decomposition,  and  its  true  molecular  weight  is  not 
known.  Salts  of  pyroarsenic  and  metarsenic  acids  may  also  be 
prepared  by  heating  secondary  and  primary  salts  of  arsenic 
acid,  but  the  acids  are  not  known  in  solution,  as  they  are  hydro- 
lyzed  by  water  at  once  to  orthoarsenic  acid.  Trisilver  arsenate, 
Ag3AsO4,  is  reddish  brown  and  insoluble ;  and  the  white,  crystal- 
line magnesium  ammonium  arsenate,  MgNH4AsO4,  is  also 
insoluble,  closely  resembling  the  corresponding  phosphate. 

Arsenic  acid  is  an  oxidizing  agent  in  concentrated  solution, 
liberating  chlorine  from  hydrochloric  acid,  while  it  is  itself 
reduced  to  arsenious  acid  or  oxide.  But  the  action  is  reversible, 


260  A  TEXTBOOK  OF  CHEMISTRY 

and  the  reverse  effect  will  occur  in  dilute  solutions,  chlorine 
oxidizing  arsenious  acid  to  arsenic  acid : 

H3As04  +  2  HC1  ^±  H3AsO3  +  C12 

In  a  neutral  or  faintly  alkaline  solution  the  equilibrium  is  so 
far  toward  the  formation  of  the  arsenate  that  the  oxidation  by 
iodine,  even,  is  practically  quantitative  and  is  used  for  the 
standardization  of  iodine  solutions  : 

Na3AsO3  + 12  +  2  NaHCO3  =  Na3AsO4  +  2  Nal  +  H2O  +  2  CO2 

Arsenic  Trichloride,  AsCl3,  may  be  prepared  by  the  direct 
union  of  arsenic  and  chlorine.  It  is  a  colorless  liquid,  which  boils 
at  130°.  It  is  almost  completely  hydrolyzed  by  water  to 
hydrochloric  acid  and  arsenious  oxide  or  acid.  Some  arsenic 
trichloride  is  still  present  in  the  solution,  however,  as  a  part 
of  the  arsenic  passes  over  on  distilling  the  solution,  while  arsenic 
does  not  escape  on  distilling  a  solution  of  arsenious  oxide. 

Sulfides  of  Arsenic.  There  are  four  sulfides  of  arsenic, 
As2S2,  As2S3,  As2$5  and  As4S3.  The  last  was  prepared  rather 
recently. 

Arsenic  Bisulfide,  or  Realgar,  As2S2,  is  found  in  nature  and 
may  be  prepared  by  melting  a  mixture  of  arsenic  and  sulfur. 
It  forms  a  red,  crystalline  mass  which  becomes  lighter  colored 
when  powdered,  and  was  formerly  used  by  artists  in  painting. 

Arsenic  Trisulfide,  or  Orpiment,  As2S3,  is  also  found  in  nature, 
and  is  prepared  artificially  by  melting  sulfur  and  arsenic  mixed 
in  the  proper  proportion.  When  prepared  in  this  way  it  forms 
a  yellow,  crystalline  mass  and  the  powder  is  used  as  a  pigment, 
especially  by  artists.  From  acid  solutions  of  arsenic  trioxide 
the  trisulfide  is  precipitated  in  an  amorphous  form.  It  is  one 
of  the  most  insoluble  sul  fides  known  and  is  scarcely  attacked 
by  the  most  concentrated  hydrochloric  acid.  It  is,  however, 
dissolved  in  the  presence  of  oxidizing  agents,  as  by  nitric  acid, 
aqua  regia  or  potassium  chlorate  and  hydrochloric  acid. 

Arsenic  Pentasulfide,  As2S5,  is  precipitated  from  a  solution  of 
arsenic  acid,  H3AsC>4,  containing  hydrochloric  acid,  apparently 
through  the  substitution  of  sulfur  for  oxygen,  giving  the  series 


ARSENIC  261 

of  acids,  H3AsSO3,  H3AsS2O2,  H3AsS3O  and  H3AsS4.  The 
last  then  dissociates  into  hydrogen  sulfide  and  arsenic  penta- 
sulfide.  (MacCay,  J.  Am.  Chem.  Soc.  24,  661  (1902) ;  Z.  anorg. 
Chem.  29,  36  (1901).  It  may  also  be  prepared  by  melting  a 
mixture  of  the  elements. 

Sulfarsenites  and  Sulfarsenates.  Arsenic  trisulfide,  As2S3, 
and  arsenic  pentasulfide,  As2Ss,  dissolve  easily  in  solutions  of 
ammonium  sulfide,  (NH4)2S,  or  sodium  sulfide,  Na2S,  giving 
solutions  of  sulfarsenites  and  sulfarsenates 

3  (NH4)2S  +  As2S3  =  2  (NH4)3AsS3 

Ammonium 
Sulfarsenite 

3  Na2S  +  As2S5  =  2  Na3AsS4 

Sodium 
Sulfarsenate 

These  compounds  may  be  considered  as  arsenites  and  arsenates 
in  which  the  oxygen  has  been  replaced  by  sulfur.  Antimony 
forms  similar  compounds,  but  bismuth  does  not  form  them  in  this 
manner  —  another  illustration  of  the  fact  that  bismuth  is  more 
distinctly  metallic  and  does  not  show  the  same  tendency  as 
arsenic  and  antimony  to  form  acid  radicals.  The  formation  of 
these  compounds  is  used  in  analytical  chemistry  to  separate 
arsenic  and  antimony  from  metals  which  are  more  metallic  in 
character  and  which  do  not  form  similar  compounds. 

From  solutions  of  the  sulfarsenites  or  sulfarsenates,  acids 
precipitate  the  arsenic  as  the  trisulfide,  As2S3,  or  the  pentasul- 
fide, As2S5. 

Colloidal  Arsenic  Trisulfide.  It  has  been  pointed  out  that 
arsenic  trisulfide  is  one  of  the  most  insoluble  compounds  known. 
It  requires  at  least  two  million  parts  of  water  to  dissolve  one  part 
of  the  sulfide.  In  spite  of  this,  however,  hydrogen  sulfide  gives 
no  precipitate  with  a  solution  of  arsenic  trioxide  in  pure  water. 
A  study  of  the  properties  of  the  solution  obtained  in  this  manner 
indicates  that  the  interaction  between  the  trioxide  and  hydrogen 
sulfide  is  practically  complete  : 

As2O3  +  3  H2S  =  As2S3  +  3  H2O 


262  A  TEXTBOOK  OF  CHEMISTRY 

The  solution  has  the  properties  of  a  typical  colloidal  "solu- 
tion "  —  a  "  solution  "  in  which  a  substance,  which  under  other 
conditions  is  insoluble  and  separates  as  a  precipitate,  remains  in 
suspension.  Such  solutions  will  pass  through  ordinary  filters  un- 
changed and  under  an  ordinary  microscope  they  appear  to  be 
homogeneous.  The  freezing  points  and  boiling  points  of  such 
solutions  are  practically  identical  with  the  freezing  point  and  boil- 
ing point  of  the  pure  solvent — in  this  case  water.  This  indicates 
that  colloids  are  not  in  the  ordinary  molecular  condition.  The 
ultramicroscope  reveals  in  many  colloidal  solutions  the  presence 
of  aggregates  which  have  a  diameter  of  from  6  to  60  //./x.1  Under 
the  influence  of  a  considerable  electrical  potential,  colloidal 
arsenic  trisulfide  moves  slowly  toward  the  anode,  indicating 
that  the  aggregates  carry  negative  charges.  In  the  case  of 
some  colloidal  solutions  the  movement  is  toward  the  cathode. 
We  may  distinguish,  therefore,  negative  colloids,  as  arsenic  tri- 
sulfide, and  positive  colloids,  as  colloidal  silver.  The  addition  of 
an  electrolyte  to  a  colloidal  solution  will  usually  cause  its  precipi- 
tation. In  general  an  electrolyte  with  bivalent  ions,  as  barium 
chloride,  BaCl2,  is  more  effective  than  one  with  univalent  ions, 
as  sodium  chloride,  but  the  effect  is  dependent  also  on  the  degree 
of  ionization  of  the  electrolyte  and  it  seems  to  be  the  cation 
(e.g.  Ba++)  which  is  effective  in  precipitating  a  negative  colloid, 
and  the  anion  (e.g.  Cl~  or  SO4  )  which  precipitates  a  positive 
colloid.  The  cation  (or  anion)  is  retained  by  the  precipitate 
and  cannot  be  washed  away,  though  it  may  be  displaced  by 
another  ion  of  the  same  sign. 

These  facts  are,  at  present,  best  understood  in  the  light  of 
the  following  theory.  In  the  colloidal  solution  aggregates  of  a 
substance  which  is  usually  insoluble  are  formed  around  negative 
or  positive  ions,  forming  in  the  first  case  negative,  in  the  second 
case  positive,  colloids.  These  aggregates  are  very  much  larger 
than  ordinary  molecules,  but  they  are  prevented  from  falling  to 

1  /*  stands  for  one  micron,  y^  of  a  millimeter.  /*/*  stands  for 
TTyW  of  a  micron  or  one  millionth  of  a  millimeter.  The  wave 
length  of  sodium  light  is  approximately  0.6  /*. 


ANTIMONY  263 

the  bottom  of  the  solution,  partly  because  they  are  still  very 
small,  but  more  because  on  account  of  their  electrical  charges 
they  are  prevented  from  cohering  with  other  similar  aggregates 
to  form  larger  particles  and  also  because  there  must  always 
be  in  the  solution,  to  balance  them  electrically,  other,  ordinary, 
ions  with  charges  of  the  opposite  signs.  If  these  aggregates 
were  to  separate  from  the  solution  as  a  precipitate,  the  solution 
would  be  electrically  positive  and  the  precipitate  negative  in  the 
case  of  arsenious  sulfide.  When  an  electrolyte,  as  barium 
chloride,  is  added  to  such  a  solution,  the  positive  barium  ions, 
Ba++,  combine  with  the  negative  ions  of  the  colloidal  arsenious 
sulfide,  forming  neutral  aggregates  which  can  then  cohere  to 
larger  aggregates  and  form  an  ordinary  precipitate.  At  the 
same  time  the  negative  chloride  ions,  Cl~,  balance  the  positive 
ions  of  the  solution,  usually  hydrogen  ions,  H+,  and  the  solution 
remains  electrically  neutral  although  the  colloid  has  separated 
from  it.  It  has  been  shown  that  in  such  a  case  the  solution 
remains  acid  in  exact  proportion  to  the  amount  of  barium  carried 
down  by  the  precipitate,  and  it  has  already  been  pointed  out 
above  that  the  barium  cannot  be  removed  from  the  latter  by 
washing. 

A  knowledge  of  the  conditions  which  govern  the  formation  and 
precipitation  of  colloids  is  often  of  very  great  importance  in 
analytical  chemistry.  The  phenomena  of  colloidal  solutions 
also  play  a  very  important  part  in  the  digestion  and  assimilation 
of  food  and  in  the  life  processes  of  both  plants  and  animals. 

Antimony,  Sb,  120.2.  Occurrence  and  Preparation.  Small 
quantities  of  antimony  are  found  free  in  nature,  but  the  element 
occurs  chiefly  in  the  mineral  stibnite,  antimony  trisulfide,  Sb2S3. 
When  this  is  heated  in  the  air,  the  sulfur  burns  away  as  sulfur 
dioxide,  SO2,  and  the  antimony  remains  as  the  tetroxide,  Sb2C>4. 
The  process  is  called  roasting  and  is  a  very  common  method  of 
treating  ores  which  contain  sulfides  of  the  metals.  The  crude 
oxide  is  then  reduced  by  heating  it  with  coke  or  charcoal  and 
suitable  substances  to  form  a  fusible  slag  with  the  impurities  of 
the  ore :  sb2O4  +  2  C  =  2  Sb  +  2  CO2 


264  A  TEXTBOOK  OF   CHEMISTRY 

Properties.  Antimony  is  a  silver-white,  brittle,  crystalline 
metal.  The  specific  gravity  is  6.52.  It  melts  at  630°  and  boils 
at  1300°. 

A  curious  form  of  the  element  known  as  explosive  antimony 
can  be  prepared  by  the  electrolysis  of  a  solution  of  antimony 
chloride,  SbCla,  in  hydrochloric  acid.  It  has  a  specific  gravity  of 
only  5.78.  When  rubbed  in  a  mortar  or  when  the  dry  sub- 
stance is  heated  to  200°,  it  explodes  violently,  with  an  appearance 
of  light  and  heat,  being  converted  into  ordinary  antimony.  The 
transformation  is  accompanied  by  the  evolution  of  about 
20  calories  per  gram. 

When  antimony  is  heated  in  the  air,  on  charcoal,  it  melts  easily, 
differing  in  this  respect  from  arsenic,  which  sublimes  without 
melting.  It  burns  slowly,  giving  vapors  of  antimony  trioxide, 
Sb2O3.  It  does  not  dissolve  in  hydrochloric  acid,  but  is  easily 
converted  into  a  mixture  of  insoluble  oxides  by  nitric  acid. 
Antimony  and  tin  are  the  only  metals  which  are  acted  upon  by 
nitric  acid  in  this  manner,  giving  insoluble  oxides  or  acids. 

Uses.  Metallic  antimony  is  a  constituent  of  many  important 
alloys,  especially  of  type  metal  (lead,  tin  and  antimony),  stereo- 
type metal  (lead,  tin,  antimony  and  bismuth),  britannia  metal 
(tin  and  antimony)  and  antifriction  metals  (lead,  antimony  and 
copper  with  a  little  bismuth)  used  for  bearings  in  machinery. 
In  type  metal  it  gives  hardness  to  the  alloy  and  also  causes  it  to 
expand  slightly  as  it  solidifies  in  the  mold,  giving  clear-cut 
type. 

Stibine,  SbH3,  is  formed  when  a  soluble  compound  of  antimony 
is  introduced  into  a  hydrogen  generator  containing  zinc  and 
hydrochloric  or  sulfuric  acid.  It  resembles  arsine  closely,  but 
gives  a  somewhat  more  sooty  spot  on  porcelain  or  in  a  glass  tube 
by  Marsh's  test.  It  is  decomposed  into  antimony  and  hydrogen 
at  a  lower  temperature  than  arsine,  the  decomposition  taking 
place  slowly  at  ordinary  temperatures,  especially  in  the  presence 
of  metallic  antimony.  For  the  methods  of  distinguishing 
between  the  deposits  of  arsenic  and  antimony,  works  on  ana- 
lytical chemistry  should  be  consulted. 


ANTIMONY  265 

Oxides  of  Antimony.  Antimony  forms  three  oxides:  anti- 
mony trioxide,  Sb4O6  (or  Sb2O3),  formed  by  burning  antimony 
in  the  air  or  by  heating  the  hydroxide,  Sb(OH)3;  antimony 
tetroxide,  Sb2O4,  formed  when  either  the  pentoxide  or  the 
trioxide  is  heated  with  free  access  of  air ;  and  antimony  pentox- 
ide, Sb2O5,  obtained  by  repeated  evaporation  of  metallic  anti- 
mony or  one  of  the  lower  oxides  with  nitric  acid  and  finally 
heating  the  residue  to  300°.  At  a  higher  temperature  it  is 
decomposed  into  the  tetroxide  and  oxygen.  The  trioxide  is  the 
only  oxide  which  can  be  converted  into  a  vapor  without  decompo- 
sition, and  so  is  the  only  one  for  which  we  really  know  the  molec- 
ular weight  and  true  formula,  Sb^e.  It  is  altogether  probable 
that  the  tetroxide  and  pentoxide  have  more  complex  formulas 
than  those  given,  and,  indeed,  it  is  quite  possible  that  the  solid 
trioxide  has  a  higher  molecular  weight  and  more  complex  formula 
than  that  of  its  vapor.  For  inorganic  compounds,  in  general, 
it  is  more  convenient  to  use  the  simplest  formulas  which  express 
the  composition  in  whole  atomic  weights.  For  this  reason  chem- 
ists continue  to  use  the  formulas  P2O3,  P2O5,  Sb2O3,  etc.,  and  the 
corresponding  names,  for  compounds  whose  molecules  are  known 
to  be  more  complex.  When  questions  of  structure  are  con- 
sidered, however,  it  is  important  to  remember  that  the  molecules 
are  more  complex  than  these  formulas  indicate. 

Antimony  hydroxide,  Sb(OH)3  or  H3SbO3,  Antimonious  Acid, 
may  be  prepared  by  the  precipitation  of  a  solution  of  tartar 
emetic  (see  below)  with  dilute  sulfuric  acid.  As  is  to  be  expected 
from  the  position  of  antimony  in  the  Periodic  System,  it  is 
amphoteric  in  character,  that  is,  both  an  acid  and  a  base.  In  a 
solution  of  a  strong  base  it  gives  up  hydrogen  and  forms  a  salt 
in  which  it  furnishes  the  acid  radical : 

H3Sb03  +  NaOH  =  NaH2SbO3.H2O 

Sodium 
Antimonite 

In  a  solution  of  a  strong  acid,  on  the  other  hand,  it  gives  up 
its  hydroxyl  and  forms  salts  in  which  the  antimony  is  the  metallic 
element :  2  Sb(OH)3  +  3  H2SO4  =  Sb2(SO4)3  +  6  H2O 


266  A  TEXTBOOK   OF  CHEMISTRY 

In  further  agreement  with  this  character,  the  hydroxide  or 
acid  loses  both  hydrogen  and  hydroxyl  (OH)  easily  even  when  in 
contact  with  water,  going  over  into  the  oxide,  Sb2O3 : 


/io-H<    JH-OJ 

Sb^-O-JHJ      JH-t-O-Sb 
\O-!H H-oi 


As  is  to  be  expected,  also,  both  classes  of  salts  are  hydrolyzed 
by  water.  The  salts  of  the  alkalies  react  strongly  alkaline  in 
solution,  while  the  salts  in  which  antimony  forms  the  metallic 
part  are  mostly  decomposed  by  water  with  the  precipitation  of  a 
basic  salt  and  liberation  of  the  free  acid.  In  these  basic  salts 


instead  of  the  group  SbO  —  H,  which  might  be  expected,  the 


group  Sb^    ,  antimonyl,  formed  from  this  by  the  loss  of  hydro- 

gen and  hydroxyl,  is  often  present  : 

Sb2(SO4)3  +  2  HOH  =  (SbO)2SO4  +  2  H2SO4 

Antimony! 
Sulfate 

Tartaric  Emetic,  KSbOC4H4O6.  One  of  the  most  interesting 
and  important  basic  salts  of  antimony  is  tartar  emetic,  or  potas- 
sium antimonyl  tartrate.  Cream  of  tartar,  or  acid  potassium 
tartrate,  KHC4H4O6,  is  the  acid  potassium  salt  of  tartaric  acid, 
H2C4H4O6,  an  acid  found  in  the  juice  of  grapes.  When  antimony 
trioxide,  Sb2O3,  is  boiled  with  a  solution  of  cream  tartar,  it  dis- 
solves, forming  tartar  emetic  : 

Sb2O3  +  2  KHC4H406  =  2  KSbOC4H4O6  +  H2O 

/o 

In  the  tartar  emetic  the  univalent  antimonyl  group,  Sbv       , 

replaces  hydrogen  .as  though  it  were  a  univalent  metal,  very 
much  as  the  ammonium  group,  NH4,  replaces  hydrogen  in  the 


ANTIMONY  267 

formation  of  ammonium  salts.  Tartar  emetic  dissolves  easily 
in  water.  It  is  sometimes  used  as  an  emetic. 

Antimonic  Acids.  Three  antimonic  acids  are  known,  cor- 
responding to  the  phosphoric  acids  of  similar  formulas  :  metanti- 
monic  acid,  HSbOa,  pyroantimonic  add,  H4Sb2C>7,  and  orthoanti- 
monic  acid,  H3SbO4.  Very  few  salts  of  the  last  are  known. 

Chlorides  of  Antimony.  Antimony  forms  three  chlorides. 
Antimony  trichloride,  SbCla,  can  be  prepared  by  dissolving  the 
trioxide,  Sb2Oa,  or  the  trisulfide,  Sb2Sa,  in  concentrated  hydro- 
chloric acid,  evaporating  the  solution  and  distilling  the  residue. 
It  is  a  solid  which  melts  at  73.2°  and  boils  at  223°.  It  is  decom- 
posed by  water  with  the  precipitation  of  the  oxychloride,  which 
may  be  called  antimony  1  chloride,  SbOCl.  From  the  method 
of  preparing  the  trichloride  it  is  evident  that  the  decomposition 
by  water  is  a  reversible  reaction,  the  direction  of  which  depends 
on  the  concentration  of  the  reacting  substances. 

*  Antimony  tetrachloride,  SbCU,  and  Hydrotetrachloroanti- 
monic  acid,  H2SbCl6.     When  solutions  of  antimony  trichloride, 
SbCls,  and  antimony  pent-achloride,  SbCls,  in  hydrochloric  acid 
are  mixed,  a  dark   brown  solution  is  formed.      The  depth   of 
color  increases  on  warming  and  decreases  on  cooling,  indicat- 
ing that  the  tetrachloride  is  formed  with  an  absorption  of  heat, 
since  its  formation  is  promoted  by  an  increase  of  temperature 

(p*  l  SbCl3  +  SbCl5  =  2  SbCl4  , 

The  tetrachloride  has  not  been  prepared  in  pure  condition,  but 
double  salts  with  other  metals  such  as  the  caesium  tetrachloro- 
antimonate,  Cs2SbCl6,  have  been  prepared.  This  corresponds 
to  an  acid,  H2SbCle,  which  probably  exists  in  the  dark  brown 
solution  referred  to  above  and  which  may  be  called  hydrotetra- 
chloroantimonic  acid. 

*  Antimony  Pentachloride,  SbCls,  is  formed  when  antimony 
is  burned  in  chlorine.     It  is  a  white  solid  at  low  temperatures, 
but  melts  at  —  6°  and  boils  at  140°.     At  the  latter  temperature 
it  dissociates,  partly,  into  the  trichloride  and  chlorine,  exactly 
as  phosphorus  pentachloride,  PCls,  does. 


268  A  TEXTBOOK  OF  CHEMISTRY 

*  Metachloroantimonic  Acid,  HSbCle.4^  H2O.  If  chlorine 
is  led  into  a  concentrated  solution  of  antimony  trichloride  in 
hydrochloric  acid  till  the  solution  becomes  colorless  or  light 
yellow  and  the  solution  is  evaporated  in  a  current  of  hydro- 
chloric acid  to  prevent  hydrolysis,  very  hygroscopic  crystals  of 
metachloroantimonic  acid  can  be  obtained.  This  may  be  con- 
sidered as  metantimonic  acid,  HSbO3,  in  which  the  three  oxygen 
atoms  have  been  replaced  by  six  chlorine  atoms.  The  freezing 
point  of  the  solution  indicates  that  the  compound  separates  into 
the  ions  H^and  SbCl~6-  The  solution  gives  a  precipitate  with 
silver  nitrate  only  after  some  time,  indicating  that  very  few 
chloride  ions,  Cl~,  are  present. 

Very  many  salts  of  this  acid  have  been  prepared,  among  which 
the  following  may  be  mentioned  : 

Potassium  metachloroantimonate    ....   KSbCle.H2O 
Calcium  metachloroantimonate       ....   Ca(SbCl6)2.9  H2O 
Aluminium  metachloroantimonate       .     .     .  Al(SbCl6)3.15  H2O 

Antimony  Trisulfide,  Sb2S3,  is  obtained  as  an  orange-red  pre- 
cipitate when  hydrogen  sulfide  is  passed  into  an  acid  solution  of  a 
salt  of  antimony.  It  also  occurs  in  nature  as  the  black  mineral, 
stibnite.  It  dissolves  readily  in  concentrated  hydrochloric  acid, 
differing  very  markedly  from  arsenic  trisulfide  in  this  regard. 

Antimony  Pentasulfide,  Sb2S5,  is  best  obtained  by  the  decom- 
position of  sodium  sulfantimonate  with  hydrochloric  acid : 

2  Na3SbS4  +  6  HC1  =  6  NaCl  +  Sb2S5  +  3  H2S 

Sulfantimonites,  M3SbS3,  and  Sulfantimonates,  M3SbS4,  may 
be  obtained  in  the  same  manner  as  the  corresponding  sulfar- 
senites  and  sulfarsenates  (p.  261).  The  alkali  salts  are  soluble 
in  water,  hence  the  sulfides  of  antimony  dissolve  in  solutions 
of  sodium  sulfide  or  ammonium  sulfide  and  may  be  separated 
in  this  way  from  the  sulfides  of  elements  which  are  very  decidedly 
metallic  in  character. 

Bismuth,  Bi,  208.  Occurrence,  Properties,  Uses.  Bismuth 
is  less  abundant  in  nature  than  arsenic  or  antimony,  as  is  usually, 
though  not  invariably,  the  case  with  elements  of  high  atomic 


BISMUTH  269 

weights.  It  is  found  mostly  in  the  free  state,  but  is  found  also 
as  the  sulfide,  Bi2S3,  both  alone  and  with  other  sulfides,  espe- 
cially with  lead  sulfide.  It  is  also  found  as  the  oxide,  Bi2O3. 
It  is  obtained  commercially  as  a  by-product  in  the  electrolytic 
refining  of  lead.  The  specific  gravity  of  the  distilled  metal  is 
9.78.  The  melting  point  is  271°.  As  bismuth  expands  on  solidi- 
fying, the  melting  point  is  lowered  by  pressure  (Principle  of 
van't  Hoff-Le  Chatelier,  p.  111).  Its  boiling  point  is  below 
1700°,  but  is  not  accurately  known. 

Bismuth  is  used  in  a  variety  of  alloys,  usually  because  it  lowers 
their  melting  points  and  renders  them  more  suitable  for  specific 
purposes.  It  is  used  in  this  way  in  stereotype  metal  to  give  an 
alloy  which  can  be  cast  in  a  papier-mache  mold  without  injuring 
it,  in  Wood's  metal  (Bi,  4  parts,  Pb,  2  parts,  Sn,  1  part,  Cd,  1  part), 
which  melts  at  60.5°,  and  is  used  for  heating  baths  in  chemical 
laboratories,  and  for  many  other  easily  fusible  alloys  used  as 
safety  plugs  in  steam  boilers  and  in  automatic  sprinklers  for 
protection  against  fire,  also  for  safety  fuses  in  electrical  work. 
The  addition  of  a  little  bismuth  has  been  found  an  advantage  in 
Babbitt  metal  and  in  other  antifriction  metals  used  for  bearings 
in  machinery. 

Oxides  of  Bismuth.  Bismuth  forms  two  well-characterized 
oxides,  BiO  and  Bi2O3.  It  also  forms  one  or  more  higher  oxides, 
called  peroxides,  for  which  we  should  expect  the  formulas  Bi2O4 
or  Bi2O5.  These  higher  oxides  do  not  seem  to  have  been  obtained 
in  a  state  of  purity,  probably  because  of  the  ease  with  which 
they  and  their  hydrates  lose  water  and  oxygen.  A  mixture  of 
these  oxides  or  hydrates  containing  some  sodium  is  prepared 
and  has  been  called,  without  good  reason,  sodium  bismuthate. 
It  oxidizes  manganese  compounds  to  permanganic  acid,  HMnC>4, 
in  nitric  acid  solutions  and  is  used  for  that  purpose  in  the  de- 
tection and  quantitative  determination  of  manganese.  Appar- 
ently no  pure  bismuthic  acid  or  salt  of  bismuthic  acid  has  been 
prepared. 

Bismuth  Chloride,  Bids,  is  formed  by  the  direct  union  of 
bismuth  and  chlorine  or  by  the  solution  of  bismuth  trioxide, 


270  A  TEXTBOOK  OF  CHEMISTRY 

Bi2O3,  in  concentrated  hydrochloric  acid.  It  melts  at  225°-230° 
and  boils  at  427°-429°.  It  dissolves  in  moderately  strong 
hydrochloric  acid  to  a  clear  solution,  but  the  addition  of  water 
causes  the  precipitation  of  bismuth  oxychloride,  or  bismuthyl 
chloride,  BiOCl,  which  is  extremely  insoluble. 

Bismuth  Nitrate,  Bi(NO3)3.5  H2O,  can  be  prepared  by  dis- 
solving either  bismuth  trioxide,  Bi2O3,  or  metallic  bismuth  in  an 
excess  of  nitric  acid  and  evaporating  to  crystallization.  When 
bismuth  nitrate  is  treated  with  water,  it  is  hydrolyzed  with  the 
formation  of  a  mixture  of  basic  nitrates,  which  varies  in  compo- 
sition according  to  the  method  by  which  it  is  prepared.  The 
simplest  of  these  compounds  are  Bi(OH)2NO3  and  BiONO3  : 

OH 

Bi  (NO,).  +  2  HOH  =  Bi—  OH  +  2  HNO3 
XNO3 


n  , 

H2O 


/.Pn  ,0 

i(-!OHl  =  Bif         + 
\LN6V  N03 


The  mixture  of  basic  nitrates  is  called  in  many  medical  works 
"  bismuth  subnitrate,"  an  antiquated  name  which  does  not 
correspond  to  modern  scientific  usage.  It  is  used  in  medicine 
and  also  as  a  slightly  antiseptic  face  powder. 

Bismuth  Trisulfide,  Bi2S3,  separates  as  a  black  or  brownish 
black  precipitate  when  hydrogen  sulfide  is  passed  into  a  solution 
of  a  soluble  bismuth  salt.  Owing  to  the  more  metallic  character 
of  bismuth,  it  does  not  dissolve  appreciably  in  solutions  of 
sodium  sulfide  or  ammonium  sulfide  as  the  sulfides  of  arsenic 
and  antimony  do.  It  dissolves  easily  in  warm  nitric  acid,  form- 
ing bismuth  nitrate,  Bi(NO3)3. 

The  following  tables  of  the  more  important  compounds  of  the 
elements  of  the  fifth  group  will  be  of  service  in  reviewing  and 
comparing  these.  Compounds  which  correspond  for  different 
elements  are  selected,  especially,  for  the  table.  Many  other 
compounds  are,  of  course,  known. 


ARSENIC,  ANTIMONY  AND  BISMUTH  271 


N,  14 

N20 

NO 

N203 

NO2,  N2O4 

N205 


NC1 


P,  31 


PC1 


PC16 


Oxides 
As,  75  Sb,  120 


AsCl 


SbCl 


Bi,  208 


Bi2O3 

PA 

As4O6                Sb4O6 

P<010 

As2O5                 Sb2O5 

Bi2O5  ? 

Chlorides 

BiCl 


H2N2O2 
HN02 


HNO3 


Acids 

H3P02 

H3P03  H3As03 

H2P03 

HP03  HAsO3 

H4P2O7  H4As2O7 

H3P04  H3As04 


H3SbO3  (Bi(OH)8) 


HSb03 


HBi03? 


Salts  of  Sulfur  Acids 


Na3PS3      Na3AsS3 
M4P2S7       Na3AsS4 


Na3SbS3 


NaBiS 


NH3 

N2H4 

N3H 


PH3 
P2H4 


Hydrides 
AsH3  SbH3 


Vanadium  (V,  51.0),  Columbium  (or  Niobium)  (Cb,  93.5), 
Tantalum  (Ta,  181.5).  These  elements,  which  are  found  in  the 
fifth  group  in  the  alternate  rows  of  the  Periodic  System,  are 


272  A  TEXTBOOK  OF  CHEMISTRY 

more  decidedly  metallic  in  their  properties,  corresponding  to 
their  positions  in  the  system ;  and  while  they  show  many  analogies 
with  the  elements  described  in  this  chapter,  they  will  be  reserved 
for  a  later  consideration  (p.  522). 

EXERCISES 

1.  What  are  the  reactions  in  Marsh's  test,  if  arsenious  oxide  is  used  ? 
What,  if  arsenic  acid  is  used  ? 

2.  If  an  arsenic  mirror  on  porcelain  is  warmed  with  ammonium 
sulfide,  it  dissolves  and  the  addition  of  hydrochloric  acid  gives  a  lemon- 
yellow  precipitate  which  does  not  dissolve  in  concentrated  hydrochloric 
acid.     Write  the  equations. 

3.  The  antimony  mirror  conducts  itself  in  a  similar  manner,  but 
gives  an  orange  precipitate  with  dilute  hydrochloric  acid  which  dissolves 
in  concentrated  hydrochloric  acid.     Write  the  equations. 

4.  What  volume  of  air  will  be  required  for  the  complete  combustion 
of  one  volume  of  arsine  ? 

5.  What  weight  of   arsenious  oxide  will  be  required  to  give  one 
pound  (453  grams)  of  Paris  green  ? 

6.  What  will  be  formed  by  the  ignition  of  magnesium  ammonium 
arsenate  ? 

7.  The  specific  gravity  of  fused  arsenic  trisulfide  is  2.76.    If  a  particle 
of  the  colloidal  sulfide  has  this  specific  gravity  and  is  one  micron  in 
diameter,  how  many  such  particles  would  there  be  in  one  gram  of  the 
sulfide?     Assuming  that  it  takes  1.5  X  1021  molecules  of  arsenious  sul- 
fide (As2S3)  to  weigh  one  gram,  how  many  molecules  would  there  be 
in  such  a  particle  of  colloidal  arsenic  trisulfide  ? 

8.  Assuming  the  formula  Bi2O5  for  bismuth  peroxide,  what  is  the 
equation  for  the  reaction  between  this  compound  and  manganese  ni- 
trate, Mg(NO3)2,  in  the  presence  of  dilute  nitric  acid  ? 


CHAPTER  XVI 
CARBON 

Carbon,  C,  12.  Occurrence.  Although  carbon  forms  only 
about  one  five-hundredth  part  of  that  portion  of  the  earth  which 
we  can  examine,  it  is  in  many  respects  the  most  important  of  all 
of  the  elements.  It  forms  an  indispensable  element  in  all  living 
organisms,  both  animal  and  vegetable,  and,  indeed,  it  seems  to  be 
the  peculiar  properties  of  carbon  rather  than  those  of  any  other 
element,  which  make  life,  as  we  know  it,  possible.  In  addition 
to  this  preeminent  role  in  living  bodies,  carbon  is  the  principal 
constituent  in  all  substances  used  for  fuel  and  is  the  element 
by  means  of  which  iron  is  reduced  from  its  ores. 

The  unique  character  of  carbon  is  suggested  by  its  position 
in  the  Periodic  System.  It  stands  in  the  first  row  midway 
between  the  most  strongly  nonmetallic  element  fluorine  and  one 
of  the  alkali  metals,  lithium.  Corresponding  to  this  position 
its  valence  is  four,  but  it  combines  both  with  the  positive  hy- 
drogen and  with  the  negative  fluorine  and  chlorine.  Even  in 
its  elementary  forms,  it  is  nonmetallic,  transparent  and  a 
nonconductor  of  electricity  in  the  diamond,  but  approaches 
the  metals  in  being  opaque  and  a  fairly  good  conductor  of  elec- 
tricity in  graphite.  This  double  character  seems  to  be  closely 
connected  with  the  power  which  carbon  atoms  have  to  combine 
with  each  other  as  well  as  with  other  elements.  The  compounds 
of  the  element  are  bewildering  in  their  variety  and  complexity. 
More  than  100,000  such  compounds  have  been  prepared  and 
analyzed,  and  some  thousands  of  new  compounds  are  discovered 
every  year.  On  account  of  their  number  and  many  peculiarities, 
which  distinguish  them  from  the  compounds  of  other  elements, 
but  also  because  of  their  importance  and  because  so  much  time 
has  been  devoted  to  their  study,  the  compounds  of  carbon  are 

273 


274  A  TEXTBOOK  OF  CHEMISTRY 

usually  considered  separately  as  a  special  subject,  called  organic 
chemistry. 

The  study  of  these  compounds  has  proved  so  important  in 
its  relation  to  the  problems  of  general  chemistry,  however,  that 
no  textbook  of  inorganic  chemistry  is  complete  without  a  descrip- 
tion of  some  of  them. 

Diamonds.  A  diamond  of  the  first  quality,  weighing  0.2  gram, 
when  properly  cut,  is  worth  approximately  $100,  while  a  kilo- 
gram of  carbon  in  the  form  of  coal  or  coke  is  worth  less  than 
one  cent.  This  fact  has  been  a  constant  challenge  to  chemists 
ever  since  the  composition  of  diamonds  was  discovered.  It  was 
not,  however,  till  near  the  close  of  the  nineteenth  century  that 
even  microscopic  diamonds  were  prepared  artificially ;  and  even 
now  the  theoretical  conditions  for  their  preparation  are  not 
fully  understood  and  no  one  has  succeeded  in  making  diamonds 
large  enough  to  be  of  commercial  value. 

In  1892  Friedel  discovered  that  a  meteor  which  had  fallen  in 
Canon  Diablo,  Texas,  contained  carbon.  By  a  careful  examina- 
tion, Moissan,  in  Paris,  demonstrated  the  presence  of  microscopic 
diamonds  in  the  material.  As  he  stated  it  afterwards,  nature 
had  been  caught  in  the  act  of  making  diamonds.  The  high 
specific  gravity  of  the  diamond  in  comparison  with  graphite  seems 
to  have  suggested  to  Moissan  that  diamonds  were  probably 
formed  under  conditions  of  high  pressure.  This  and  the  dis- 
covery of  diamonds  in  the  meteorite  suggested  to  him  the  follow- 
ing experiment.  A  mass  of  iron  was  heated  to  a  very  high  tem- 
perature, 3000°-3500°,  in  an  electric  furnace  and  some  pure 
sugar  charcoal  was  dissolved  in  the  hot  iron.  The  mass  was 
then  suddenly  thrust  into  water.  This  caused  the  exterior 
surface  to  solidify  while  the  interior  was  still  fluid  and  intensely 
hot.  As  the  interior  cooled,  some  of  the  dissolved  carbon  sepa- 
rated in  crystalline  form,  and,  as  iron  containing  carbon  ex- 
pands on  solidifying,  the  interior  portions  were  subjected  to  an 
enormous  pressure.  The  lowering  of  the  melting  point  of  the 
iron  by  the  pressure  may  also  have  had  something  to  do  with  the 
success  of  the  experiment.  When  cold,  the  iron  was  dissolved 


DIAMONDS  275 

in  acids,  and  silicon,  graphite  and  other  substances  were  removed 
by  oxidation,  solution,  and  finally  by  treatment  with  a  liquid 
having  a  specific  gravity  greater  than  that  of  graphite  and  less 
than  that  of  the  diamond.  A  few  minute  crystals  were  dis- 
covered which  were  heavier  than  this  liquid.  By  carefully 
rubbing  one  of  the  crystals  against  the  surface  of  a  ruby,  it  was 
shown  that  the  surface  of  the  latter  was  scratched.  Only  the 
diamond  and  carborundum  are  known  to  be  harder  than  the 
ruby,  and  the  specific  gravity  of  carborundum  is  less  than  that 
of  the  artificial  diamonds.  After  securing  enough  of  the  crystals 
to  weigh  a  few  milligrams,  they  were  burned  in  a  current  of 
oxygen  and  the  carbon  dioxide  formed  was  absorbed  and  weighed. 
Twelve  parts  by  weight  of  the  crystals  gave  44  parts  by  weight 
of  carbon  dioxide.  The  crystals  consisted,  therefore,  of  pure 
carbon  and  were  in  reality  diamonds. 

Until  the  beginning  of  the  eighteenth  century  diamonds  had 
been  found  only  in  India.  In  1727  they  were  discovered  in 
Brazil,  at  the  beginning  of  the  nineteenth  century  in  the  Ural 
mountains,  and  in  1867  at  Kimberley  in  South  Africa.  The 
Kimberley  mines  now  furnish  most  of  the  diamonds  for  the 
world's  market.  The  product  of  the  mines  is  valued  at  about 
$15,000,000  annually. 

The  diamond  crystallizes  in  octahedra  and  cubes  of  the  iso- 
metric system.  Its  specific  gravity  is  3.5.  It  is  the  hardest 
known  substance,  and  is  not  dissolved  or  oxidized  by  any  known 
liquid  or  gas  at  ordinary  temperatures.  Its  index  of  refraction 
is  extraordinarily  high,  being  2.417  for  sodium  light.  This  and 
the  high  dispersive  power  give  to  diamonds,  which  are  cut  so 
as  to  accentuate  these  properties,  the  ability  to  reflect  and  re- 
fract light  in  such  a  manner  as  to  produce  brilliant  colors. 

For  cutting  glass,  an  edge  produced  by  cleavage  must  be  used. 
For  diamond  drills  to  be  used  in  boring  in  rocks  in  such  a  manner 
that  a  solid  core  can  be  removed,  inferior  black  diamonds,  called 
carbonado,  set  into  the  ends  of  tubes,  are  employed.- 

When  heated  to  a  high  temperature  with  exclusion  of  air,  the 
diamond  is  changed  to  graphite.  When  heated  in  oxygen,  carbon 


276  A  TEXTBOOK  OF  CHEMISTRY 

dioxide  begins  to  be  formed  at  720°  and  the  diamond  takes  fire 
and  burns  at  800°-850°. 

Graphite,  a  second  crystalline  form  of  carbon,  is  found  in  nature, 
especially  in  Ceylon,  Siberia,  England  and  Canada.  It  may  be 
prepared  artificially  by  crystallizing  carbon  from  cast  iron, 
from  one  to  two  per  cent  of  graphite  being  left  behind  when  gray 
cast  iron  is  dissolved  in  acids.  It  is  also  formed  when  any  form 
of  carbon  is  heated  to  a  very  high  temperature  in  an  electric 
furnace.  The  presence  of  silicon,  aluminium,  iron  and  other 
elements  seems  to  assist  in  the  transformation  by  the  interme- 
diate formation  of  carbides.  In  the  Acheson  process,  which  has 
acquired  considerable  technical  importance,  impure  carbon,  as 
coke  or  anthracite,  is  found  more  suitable  than  pure  carbon. 

Graphite  crystallizes  in  six-sided  leaflets  of  the  monoclinic 
system.  It  is  soft  and  has  a  gray,  metallic  luster  and  gives  a 
metallic  streak.  Its  specific  gravity  when  pure  is  2.255.  In 
oxygen  it  begins  to  give  carbon  dioxide  at  570°  and  takes  fire 
at  690°.  At  very  high  temperatures  it  seems  to  be  the  most 
stable  form  of  carbon,  into  which  all  other  forms  tend  to  pass. 
The  total  energy  of  graphite  at  ordinary  temperatures  is,  how- 
ever, greater  than  that  of  the  diamond  at  ordinary  temperatures, 
as  is  evident  from  the  following  table  of  heats  of  combustion,  as 
determined  by  Berthelot : 

12  grams  of  diamond  give  94,310  small  calories. 
12  grams  of  graphite  give  94,810  small  calories. 
12  grams  of  amorphous  carbon  give  97,650  small  calories. 

From  this  table  it  appears  that  if  12  grams  of  graphite  could 
be  changed  to  diamond  at  ordinary  temperatures,  500  small 
calories  would  be  evolved. 

Graphite  is  used  for  "  lead  "  pencils,  as  a  lubricant,  especially  for 
surfaces  of  wood  and  where  bearings  are  subjected  to  a  high  tem- 
perature, and  in  making  crucibles  for  use  at  very  high  tempera- 
tures for  melting  and  casting  steel  and  difficultly  fusible  alloys. 
For  the  last  purpose  it  must  be  mixed  with  some  fire  clay, 
which  binds  the  particles  of  graphite  together  and  also  protects 


CARBON  277 

the  graphite  from  burning  by  giving  an  incombustible  surface. 
Graphite  is  also  used  as  a  lubricant  in  a  colloidal  solution  in 
oil  (Acheson).  It  is  used  for  stove  polish,  to  protect  iron  from 
rusting. 

Amorphous  Carbon.  When  almost  any  compound  of  carbon  is 
heated,  it  will  decompose  with  the  separation  of  charcoal  or  car- 
bon. It  is,  however,  extremely  difficult  to  obtain  perfectly  pure 
carbon  in  this  manner.  The  purest  carbon  is  obtained  by  heat- 
ing sugar,  which  contains  only  carbon,  hydrogen  and  oxygen,  to  a 
high  temperature,  but  it  seems  doubtful  whether  the  last  traces  of 
hydrogen  can  be  expelled  without  the  use  of  a  temperature  which 
would  convert  the  amorphous  carbon  partly  into  graphite.  The 
hydrogen  may  be  almost  completely  removed,  however,  by 
heating  the  charcoal  in  a  current  of  chlorine  at  1000°.  There 
seems  to  be  no  form  of  amorphous  carbon  which  can  be  properly 
spoken  of  as  a  definite  chemical  individual,  as  the  density  and  kin- 
dling temperatures  and  conductivity  for  electricity  vary  gradually 
from  a  form  which  has  a  density  of  1.45  and  a  kindling  tempera- 
ture of  300°,  up  to  forms  which  approach  closely  to  graphite  in 
their  properties.  This  fact  is  doubtless  intimately  connected 
with  the  almost  infinite  variety  of  ways  in  which  carbon  atoms 
unite  with  each  other  in  the  compounds  of  carbon. 

A  great  variety  of  impure  forms  of  amorphous  carbon  are 
known.  All  of  these  contain  at  least  some  hydrogen  and  most 
of  them  contain  oxygen  and  other  elements. 

Lampblack  is  the  soot  deposited  from  substances  rich  in  carbon, 
such  as  naphthalene,  rosin,  petroleum,  etc.,  burning  with  a  smoky 
flame.  It  is  used  as  a  pigment,  especially  in  printers'  ink.  The 
insoluble  and  indelible  quality  of  carbon  makes  such  ink  even 
more  permanent  than  the  paper  on  which  it  is  printed. 

Wood  Charcoal  is  manufactured  by  piling  wood  in  heaps, 
covering  it  with  sod  and  setting  fire  to  it  in  such  a  manner  that 
only  a  portion  of  the  wood  burns  while  the  heat  converts  the  re- 
mainder into  charcoal.  The  process  is  wasteful  and  has  been 
largely  replaced  by  methods  of  charring  in  retorts  or  chambers 
so  arranged  that  the  wood  tar,  wood  alcohol  and  acetic  acid, 


278  A  TEXTBOOK  OF  CHEMISTRY 

which  are  formed  by  the  decomposition  of  the  wood,  may  be 
recovered.  Combustible  gases,  which  are  also  formed,  are  util- 
ized to  heat  the  retorts  or  chambers.  Charcoal  is  used  by  tinners, 
in  small  charcoal  furnaces,  for  filtering  alcohol  to  purify  it  and 
for  the  manufacture  of  a  high  grade  of  iron.  It  was  formerly 
used  in  very  large  quantities  for  this  last  purpose,  but  has  been 
almost  entirely  displaced  by  coke  and  coal. 

Charcoal,  because  of  the  infusible  character  of  all  forms  of  car- 
bon, retains  the  original  structure  of  the  wood  and  contains  an 
immense  number  of  pores  of  microscopic  size.  Apparently  for 
this  reason  freshly  ignited  charcoal  will  absorb  many  times  its 
volume  of  gases,  especially  of  those  gases  which  are  easily  lique- 
fied, such  as  ammonia,  or  hydrogen  sulfide.  This  phenomenon 
is  called  adsorption,  and  seems  to  be  due  to  the  condensation  of 
the  gas  on  the  very  large  surface  offered  by  the  porous  charcoal. 
Charcoal  cooled  by  liquid  air  is  a  very  efficient  means  for  ab- 
sorbing residual  gases  and  producing  a  high  vacuum. 

Charcoal  was  formerly  much  used  in  domestic  water  filters, 
but  it  has  been  found  that  such  filters  are  only  very  temporarily 
effective. 

Animal  Charcoal  and  Bone  Black  are  obtained  by  charring 
the  refuse  of  slaughterhouses  and  bones.  They  show  the  prop- 
erty of  absorption,  especially  for  coloring  matters,  in  a  very 
high  degree,  and  are  used  in  the  removal  of  color  from  sirups 
and  other  liquids,  as  in  the  purification  of  sugar.  They  are  also 
frequently  used  for  the  purification  of  organic  compounds  in 
chemical  laboratories.  For  such  use  they  should  be  purified  by 
treatment  with  acids  to  remove  calcium  phosphate  and  other 
mineral  matters  which  they  contain. 

Coke  bears  very  much  the  same  relation  to  bituminous  coal 
that  charcoal  does  to  wood.  It  is  still  manufactured  in  America, 
chiefly  in  the  so-called  "  beehive "  ovens  —  hemispherical 
chambers  built  of  brick,  12  feet  in  diameter  and  seven  and  one- 
half  feet  high.  These  are  charged  with  coal  while  still  hot  from 
a  previous  charge,  and  the  volatile  matter  given  off  from  the 
coal  takes  fire  and  burns  within  the  oven,  over  the  surface  of  the 


CARBON  279 

coal,  furnishing  the  heat  necessary  to  convert  the  coal  into  coke. 
An  opening  on  one  side  of  the  oven  supplies  air  and  the  products 
of  combustion  escape  through  a  circular  opening  at  the  top. 
This  wasteful  method  is  being  slowly  replaced  by  methods  of 
coking  in  retorts,  by  the  use  of  which  it  is  possible  to  recover  the 
tar,  ammonia  and  combustible  gases.  The  gas  produced  is 
formed  in  excess  of  what  is  necessary  for  heating  the  retorts, 
and  may  be  used  in  part  for  other  purposes,  while  in  the  beehive 
ovens  nothing  is  saved  except  the  coke. 

Coke  is  used  chiefly  in  blast  furnaces  for  the  production  of  cast 
iron.  It  is  used  in  some  other  metallurgical  processes  and  to  a 
limited  extent  as  a  domestic  fuel. 

Gas  Carbon.  Carbon  Electrodes.  On  the  walls  of  the  retorts 
used  in  the  manufacture  of  illuminating  gas,  carbon  is  deposited 
from  the  decomposition  of  carbon  compounds  in  the  volatile 
matter  given  off  by  the  coal.  As  a  result  of  the  prolonged 
heating  it  assumes  a  semicrystalline,  dense  form,  approaching 
graphite  in  its  properties.  It  has  a  density  of  1.9-2.0,  a  very 
high  kindling  temperature,  and  is  a  fairly  good  conductor  of 
electricity.  This  gas  carbon,  or  frequently,  also,  anthracite, 
petroleum  coke,  or  some  other  form  of  amorphous  carbon,  is 
mixed  with  a  little  coal  tar  or  some  petroleum  product  for  a 
binding  material  and  molded  into  various  forms  for  use  as 
electrodes  in  the  electroylsis  of  sodium  chloride,  aluminium  com- 
pounds or  other  substances,  for  use  in  electrical  furnaces  and  for 
the  carbon  electrodes  of  arc  lights.  The  mixture  is  subjected 
to  a  hydraulic  pressure  of  500  atmospheres  to  render  it  as  dense 
as  possible,  and  is  then  heated  to  a  temperature  of  1200°-! 400° 
for  24-48  hours,  till  all  volatile  compounds  have  been  expelled 
and  the  carbon  has  become  dense  and  hard  and  a  good  electrical 
conductor.  Such  electrodes  are  scarcely  attacked  by  the  chlo- 
rine evolved  in  the  electrolysis  of  a  solution  of  sodium  chloride 
and  are  scarcely  affected  when  heated,  out  of  contact  with  air,  to 
any  temperature  below  that  of  the  electric  arc.  Their  resistance 
to  oxidizing  agents  in  electroylsis  maybe  further  increased  by  con- 
verting them  partly  or  wholly  into  graphite  in  an  electric  furnace. 


280 


A  TEXTBOOK  OF  CHEMISTRY 


Coal.  Very  much  the  same  process  which  occurs  rapidly 
when  wood  is  heated  seems  to  have  gone  on  slowly  through  some 
hundreds  of  thousands  or  millions  of  years  with  vast  quantities 
of  woody  material  accumulated  during  certain  periods  of  geolog- 
ical time  and  afterwards  covered  with  thick  layers  of  clay  and 
other  materials,  beneath  the  surface  of  ancient  oceans.  Wood 
consists  chiefly  of  carbon,  hydrogen,  oxygen  and  nitrogen,  with 
small  quantities  of  mineral  matter.  The  transformation  to 
coal  has  been  occasioned  by  the  gradual  loss  of  oxygen  and  some 
of  the  hydrogen,  in  such  a  manner  that  the  per  cent  of  carbon 
gradually  increases.  The  per  cent  of  the  oxygen  decreases  until 
that  element  nearly  disappears  in  anthracite.  The  per  cent  of 
hydrogen  decreases  only  slightly  till  the  last  stage  —  the  trans- 
formation to  anthracite  —  is  reached.  It  seems  probable  that 
this  last  transformation  occurred  at  a  more  elevated  temperature. 
The  changes  in  composition  which  have  taken  place  during  these 
transformations  are  apparent  in  the  following  table: 


CHANGES  OF  WOOD  MATERIAL  DURING  GEOLOGICAL  TIME 


PERCENTAGE  COMPOSITION 
EXCLUSIVE  OP  MOISTURE 

PERCENT- 

PERCENT- 

CALORIFIC 
VALUE  ; 

MATERIAL 

AND  ASH 

AGE  OF 

AGE  OF 

CALORIES 

Car- 

Hydro- 

Oxy- 

Nitro- 

ASH 

MOISTURE 

PER  KIL- 
OGRAM 

bon 

gen 

gen 

gen 

Wood  —  Oak    . 

50.35 

6.04 

43.52 

0.09 

0.37 

20.00  2 

3696 

Peat     .... 

59.70 

5.70 

33.04 

1.56 

11.84 

14.24  2 

3979 

Brown  Lignite, 

North  Dakota 

74.88 

4.99 

19.12 

1.01 

9.35 

35.38 

3846 

Black     Lignite, 

Colorado   .     . 

76.83 

5.34 

16.29 

1.54 

5.99 

18.68 

5635 

Bituminous, 

Illinois      .     . 

83.42 

5.29 

9.52 

1.77 

11.28 

8.50 

6542 

Semibitumin- 

ous,  West  Vir- 

ginia      Poca- 

hontas   .    .    . 

91.50 

4.38 

3.07 

1.05 

6.55 

3.67 

7939 

Anthracite    .    . 

93.76 

2.72 

3.11 

0.41 

10.80 

2.18 

7216 

Charcoal  .    .    . 

84.11 

1.53 

14.36 

2.50 

6626 

Coke     .... 

95.47 

0.67 

2.82 

1.04 

14.80 



6768 

1  Table  prepared  by  Professor  S.  W.  Parr. 


2  Air  dry. 


CARBON  281 

The  properties  of  the  coals  of  different  kinds  follow  from  their 
composition.  Peat,  lignites,  and  bituminous  coals  increase 
progressively  in  calorific  value  as  the  amounts  of  moisture  and 
oxygen  decrease.  The  oxygen  in  these  coals  may  be  consid- 
ered as  combined  with  either  carbon  or  hydrogen  and  lessens  by 
so  much  the  amount  of  these  elements  which  can  evolve  heat  by 
combustion.  Bituminous  coals  may  equal  or  even  exceed  an- 
thracite coals  in  calorific  value  because  a  pound  of  hydrogen 
gives  by  its  combustion  more  than  three  times  as  much  heat  as  a 
pound  of  carbon.  Such  coals,  however,  give  off  volatile  products 
which  burn  with  a  smoky  flame,  and  hence  require  much  greater 
care  in  use  to  secure  effective  combustion. 

Three  classes  of  bituminous  coals  are  distinguished :  coking 
coals,  which  sinter  together  when  heated,  giving  a  hard,  coherent 
coke ;  noncoking  coals,  which  do  not  sinter,  or  sinter  imperfectly, 
giving  a  friable  coke ;  and  cannel  coals,  coals  of  a  peculiar,  homo- 
geneous structure  and  conchoida.l  fracture,  which  burn  with  a 
brilliant  flame  like  that  of  a  candle.  These  last  coals  are  used 
in  the  manufacture  of  illuminating  gas.  The  difference  between 
coking  and  noncoking  coals  seems  to  be  occasioned  by  the 
presence  or  absence  of  some  compound  whose  character  is  little 
understood  and  which  does  not  seem  to  be  closely  connected  with 
the  percentage  composition  of  the  coal. 

Chemical  Properties  of  Carbon.  The  most  remarkable  prop- 
erty of  carbon  is  the  extreme  slowness  with  which  it  reacts  at 
ordinary  temperatures  with  elements  for  which  it  has  a  very 
strong  affinity  at  high  temperatures.  This  is  especially  true  in 
its  relation  to  oxygen.  Elementary  carbon  in  either  of  its  three 
forms  may  remain  in  contact  with  air  for  centuries  without  any 
apparent  effect,  although  at  very  high  temperatures  there  seems 
to  be  almost  no  element  from  which  carbon  will  not  take  away 
oxygen.  Practically  all  organic  compounds  must  be  considered 
as  in  a  state  of  unstable  equilibrium  in  the  presence  of  oxygen, 
for  we  have  only  to  heat  them  to.  their  kindling  temperature 
when  they  will  burn  with  very  considerable  evolution  of  heat. 
On  this  property  depends  the  use  of  carbon  and  its  compounds 


282  A  TEXTBOOK  OF  CHEMISTRY 

for  fuel,  for  the  reduction  of  iron  ores  and  for  other  metallurgical 
operations.  On  this  property,  too,  depends  the  existence  of  the 
almost  infinite  variety  of  compounds  which  form  the  material 
basis  of  the  world  of  life  —  compounds  showing  all  possible 
gradations  in  their  content  of  energy  and  the  existence  of  which 
would  be  impossible  if  carbon  passed  quickly,  as  most  other 
elements  do,  to  the  most  stable  forms  of  combination. 

In  combination  with  other  elements  carbon  is  almost  always 
quadrivalent.  Methane,  CH4,  and  carbon  dioxide,  CO2,  may  be 
considered  as  the  most  typical  compounds.  Carbon  is  bivalent 
in  only  a  very  few  compounds  and  exclusively  in  combination 
with  atoms  or  groups  of  a  negative  character,  as  in  carbon 
monoxide,  C=O,  hydrocyanic  acid,  H — N=C,  and  fulminic  acid, 
H — O — N=C.  There  is  some  evidence  that  it  may  be  tri- 
valent  in  very  unusual  combinations,  but  it  is  then  extraordi- 
narily reactive. 


CHAPTER  XVII 


HYDROCARBONS. 


ILLUMINATING   AND    PRODUCER    GAS. 
FLAME 


CARBON  combines  with  hydrogen  to  form  many  hundreds  of 
compounds,  called  hydrocarbons.  These  compounds  may  be 
classified  in  a  number  of  series  in  accordance  with  their  com- 
position. The  following  table  illustrates  the  relations  which 
have  been  found  between  the  formulas  of  successive  hydrocar- 
bons in  any  series  and  between  the  hydrocarbons  of  different 
series.  Each  series  is  named  from  its  first  member.  In  the 
table  only  one  series  for  a  given  general  formula  is  given,  but,  as 
will  be  seen  below,  for  each  general  formula,  except  the  first, 
two  or  more  series  are  possible.  The  series  of  the  formula 
CwH2n_4  is  omitted  from  the  table  because  the  lower  members 
of  this  series  are  relatively  unimportant. 


MARSH  GAS  SERIES 
CwH2n+2 

ETHYLENE  SERIES 
CwH2n 

ACETYLENE  SERIES 
CnH2n-2 

BENZENE  SERIES 
C«H2w-6 

Methane     CH4 



L- 

Ethane       C2H6 

Ethene       C2H4 

Acetylene  C2H2 



Propane      C3H8 

Propene     C3H6 

Propine     C3H4 



Butane        C4Hio 

Butene       C4H8 

Butine       C4H6 



Pentane      C5Hi2 

Pentene     C5Hi0 

Pentine      CsHg 



Hexane       C6Hi4 

Hexene      C6H12 

Hexine       C6Hi0 

Benzene   C6H6 

Heptane      C7H16 

Heptene     C7Hi4 

Heptine     C7Hi2 

Toluene    CyHg 

Octane        CgHig 

Octene       C8Hi6 

Octine        C8Hi4 

Xylene     C8Hi0 

The  existence  of  these  compounds  may  be  explained  very 
simply  on  the  hypothesis  that  carbon  is  quadrivalent  and  that 
carbon  atoms  unite  readily  with  each  other.  This  hypothesis 

283 


284  A  TEXTBOOK  OF  CHEMISTRY 

gives  us  the  following  formulas  for  the  first  three  members  of  the 
Marsh  gas  series : 

H      H      H 

I         I        I 
H— C— C— C— H 

I        I        I 
H      H      H 

Propane 

For  the  fourth  member  of  the  series  the  theory  suggests  two 
formulas : 


H 

i 

H      H 

i        i 

H  —  C  —  H 

i 

1        1 

H—  C—  C- 

1        | 

H 

1 
H 

Methane 

1        1 
H      H 

Ethane 

H    H    H 
H— C— C— C— H 


H    H    H    H  H 

Normal  Butane  H — C — H 

Boiling  point,  +1°  | 

H 

Isobutane 
Boiling  point,  -11.5° 

These  two  hydrocarbons  have  been  prepared  by  methods 
which  leave  no  doubt  as  to  the  structure  of  each. 

For  the  series  CTCH2/l  the  theory  suggests  that  we  may  have 
compounds  in  which  carbon  atoms  are  doubly  united  and  also 
compounds  in  which  there  is  a  ring  of  carbon  atoms.  Thus  we 
may  have : 

H    H    H  H\XH 

III  C 

H— C— C  =  C— H  and  Hv    X\    /H 


H  H/  H 


Propylene  (Propene)  Cyclopropane 

Boiling  point,  -37°  Boiling  point,  -35° 

Both  of  these  compounds  are  known,  and  the  structure  has 
been  established  by  a  study  of  the  methods  of  preparation  and 
of  their  conduct  toward  various  reagents. 


HYDROCARBONS  285 

For  the  series  CnH2n_2  there  are  four  possibilities  :  one  triple 
union,  as  in  acetylene,  H  —  C  =  C  —  H  ;  two  double  unions,  as 
in  butadiene,  CH2=CH  —  CH=CH2;  cyclic  compounds  with 

/CH2—  CH 
one  double  union,  as  cylopentene,  CH  2  II    ;  and  compounds 

\CH2—  CH 
with  two  cycles,  or  rings,  as  dekahydronaphthalene, 


CH2—  CH2—  CH—  CH2—  CH2 

The  illustrations  given  would  seem  to  include  all  of  the  types 
of  combination  possible  for  carbon  and  hydrogen  atoms,  since 
quadruple  unions  between  carbon  atoms  would  be  impossible 
for  atoms  which  are  united  to  any  other  atoms.  There  are, 
however,  certain  other  relations  which  seem  to  depend  on  the 
arrangement  of  the  atoms  in  space. 

Some  of  the  combinations  of  these  forms  give  properties  which 
would  not  be  expected  from  the  formulas  of  the  compounds. 
This  is  especially  true  of  the  benzene  series,  in  all  of  the  com- 
pounds of  which  there  is  a  ring  of  six  carbon  atoms,  each  of  which 
is  united  to  one  hydrogen  atom  or  to  some  other  univalent  atom 
or  group.  The  simplest  formula  of  benzene  is  that  proposed 
byKekule: 

H 


H—  C/      ^C—  H 

II  I 

H  —  C         jj  —  H 


H 

and  many  of  the  properties  and  reactions  of  the  hydrocarbon  are 
satisfactorily  represented  by  this  formula,  but  for  other  proper- 
ties it  does  not  give  a  satisfactory  account.  A  further  considera- 
tion of  this  and  similar  questions  is  impossible  here. 


286  A  TEXTBOOK  OF  CHEMISTRY 

Marsh  Gas  or  Methane,  CH4.  When  decaying  leaves  in  the 
bottom  of  a  pond  are  stirred,  bubbles  of  a  combustible  gas  con- 
sisting largely  of  marsh  gas  or  methane,  CH4,  rise  to  the  surface. 
The  same  gas  escapes  from  seams  of  coal,  doubtless  having  been 
formed  in  a  similar  manner.  In  coal  mines  it  is  called  fire  damp. 
A  combustible  gas  consisting  very  largely  of  methane  is  often 
found  stored  in  large  quantities  in  porous  sandstones  or  lime- 
stones lying  beneath  an  impervious  layer  of  shale  so  situated  as 
to  form  a  large  inverted  reservoir.  The  gas  is  usually  under 
strong  hydrostatic  pressure  from  water  beneath.  When  such  a 
reservoir  is  pierced  by  boring  from  above,  the  gas  escapes  through 
the  opening,  and  is  known  as  natural  gas.  When  almost  any 
kind  of  organic  matter  is  heated,  methane  is  one  of  the  products 
of  decomposition,  hence  it  is  always  a  constituent  of  illuminat- 
ing gas  made  by  heating  coal,  oil  or  wood.  At  a  white  heat 
carbon  will  unite  directly  with  hydrogen  to  form  methane : 

C  +  2  H2  ^±  CH4 

The  equilibrium  of  the  reaction  is,  however,  very  far  on  the 
side  toward  the  decomposition  of  methane  into  carbon  and  hy- 
drogen. Water  gas  (p.  296)  usually  contains  a  very  small 
amount  of  methane,  which  is  probably  formed  by  the  direct 
union  of  the  elements. 

In  the  laboratory  methane  is  most  easily  prepared  on  a  small 
scale  by  heating  a  mixture  of  sodium  acetate,  NaC2H3O2,  and 
soda  lime,  which  is  a  mixture  of  sodium  hydroxide,  NaOH,  and 
slaked  lime,  Ca(OH)2.  The  slaked  lime  is  added  to  render  the 
mixture  infusible : 

NaC2H3O2  +  NaOH  =  Na2CO3  +  CH4 

Sodium 
Carbonate 

Methane  is  the  lightest  gaseous  compound  known.  It  may  be 
condensed  to  a  liquid,  which  boils  at  —  164°.  It  is  a  compara- 
tively stable  compound,  and  its  kindling  temperature  is  higher 
than  that  of  hydrogen  or  than  that  of  most  other  hydrocarbons. 


SUBSTITUTION  287 

Mixtures  of  the  gas  with  oxygen  or  with  air  explode  violently 
when  ignited.  It  burns  from  a  jet  with  a  blue  flame,  which 
gives  very  little  light. 

Substitution.  When  a  mixture  of  methane  and  chlorine  is 
exposed  to  the  sunlight,  a  double  decomposition  occurs  in  which 
one  atom  of  the  chlorine  molecule  combines  with  an  atom  of 
hydrogen  while  the  other  combines  with  the  carbon,  apparently 
taking  the  place  of  the  hydrogen.  This  process,  which  occurs 
in  a  great  variety  of  reactions  of  organic  compounds,  is  called 
substitution : 

H  H 

H— C— H  +  Cl— Cl      =      H— C— Cl  +  H— Cl 

H 

Methyl  Chloride 

The  process  may  be  continued  till  all  of  the  hydrogen  has  been 
replaced  by  chlorine : 

CH3C1  +  C12  =  HC1  +  CH2C12 

Methylene 
Chloride 

CH2C12  +  C12  =  HC1  +  CHC13 

Chloroform 

CHC13  +  Cla  =  HC1  +  CC14 

Carbon 
Tetrachloride 

Practically,  a  mixture  of  the  four  products  is  obtained  by  this 
process,  so  that  these  reactions  have  only  a  theoretical  interest. 

The  Davy  Safety  Lamp.  The  explosive  character  of  mixtures 
of  methane  and  air  has  been  mentioned.  Early  in  the  nineteenth 
century  the  frequent  explosions  of  fire  damp  in  coal  mines  in  Eng- 
land, often  causing  the  death  of  miners,  led  to  a  request  of  Sir  Hum- 
phrey Davy  that  he  should  investigate  the  matter  and  endeavor 
to  suggest  a  remedy.  He  found,  as  the  result  of  his  investigation, 
that  mixtures  containing  one  volume  of  methane  with  more  than 
six  and  less  than  fourteen  volumes  of  air  would  explode  when 


288 


A  TEXTBOOK  OF  CHEMISTRY 


Fig.  80 


ignited.1  Outside  of  these  limits  explosions  do  not  so  readily 
occur.  He  also  found  that  the  kindling  temperature  of  such 
mixtures  is  comparatively  high,  requir- 
ing contact  with  a  surface  heated  nearly 
or  quite  to  dull  redness.2  This  prop- 
erty can  be  easily  illustrated  for  illu- 
minating gas  by  pressing  a  cold  piece  of 
wire  gauze  down  over  the  flame  of  a 
Bunsen  burner  (Fig.  80).  Unburned 
gas  from  the  center  of  the  flame  will 
pass  through  the  gauze  and  the  mixture 
of  gas  and  air  above  the  gauze  will  not 
take  fire  until  the  latter  becomes  nearly 
red-hot.  On  the  basis 
of  this  fact  Sir  Hum- 
phrey Davy  invented  the  Davy  Safety  Lamp, 
which  has  the  flame  of  the  lamp  completely 
surrounded  by  wire  gauze  (Fig.  81).  The 
lamp  must,  of  course,  be  lighted  and  closed 
before  the  miner  enters  the  mine.  The  lamps 
are  usually  so  constructed  that  they  can  be 
locked.  Sometimes  a  lock  is  used  which  can 
only  be  opened  with  a  strong  electromagnet, 
so  that  it  will  be  impossible  for  the  miner  to 
open  the  lamp  in  the  mine.  While  the  danger 
of  explosions  is  greatly  lessened,  it  is  not  en- 
tirely removed  by  the  use  of  the  lamp.  If 
much  fire  damp  is  present,  a  cap  of  flame  ap- 
pears inside,  above  the  flame  of  the  lamp ;  and  this  might,  some- 
times, heat  the  wire  gauze  to  the  kindling  temperature  of  the 
mixture.  In  blasting,  too,  the  sudden  vibration  from  the  blast 

1  Later  investigations  have  shown  that  these  results  are  only  a 
very  rough  approximation.     One  authority  states  that  air  contain- 
ing 2  per  cent  of  methane  may  be  dangerous. 

2  V.  Meyer,  many  years  later,  found  that  the  kindling  temperature 
of  mixtures  of  methane  and  oxygen  in  glass  vessels  is  650°-680°. 
Ber.  26,  2429. 


Fig.  81 


PETROLEUM  289 

may  carry  the  flame  through  so  quickly  that  it  is  not  cooled 
below  the  kindling  temperature  by  the  gauze. 

Mixtures  of  very  fine  dust  containing  organic  matter  with  air 
may  explode  in  the  same  manner  as  fire  damp.  In  this  way 
destructive  explosions  have  occurred  with  coal  dust  in  dry  mines, 
with  flour  in  flour  mills,  and  with  similar  dust  in  other  factories. 

Homologues  of  Methane.  On  examining  the  formulas  of  the 
hydrocarbons  given  in  the  table  it  will  be  found  that  the  succes- 
sive members  of  any  series  differ  by  one  carbon  and  two  hydrogen 
atoms.  The  reason  for  this  is  apparent  from  the  structural 
formulas  of  methane, .  ethane,  propane,  etc.  Any  hydrocarbon 
in  such  a  series  is  called  a  homologue  of  the  lower  members  of 
the  series  and  the  series  is  called  a  homologous  series. 

Petroleum  is  found  in  great  underground  reservoirs  somewhat 
similar  to  those  containing  natural  gas.  Oil  fields  have  been 
found  widely  distributed  in  America,  especially  in  Pennsylvania, 
Ohio,  Indiana,  Illinois,  Kansas,  Texas,  California  and  Canada. 
A  large  field  is  found  in  the  Caucasus,  and  doubtless  very  many 
undiscovered  fields  exist  in  other  parts  of  the  world.  Petroleum 
consists  chiefly  of  a  very  complex  mixture  of  hydrocarbons. 
The  petroleum  from  different  localities  differs  very  considerably 
in  the  nature  of  the  hydrocarbons  which  it  contains  and  also  in 
the  amount  of  the  compounds  of  sulfur  which  are  present.  The 
Pennsylvania  petroleum  consists  largely  of  homologues  of  me- 
thane. California  and  Caucasus  petroleum  contain  compounds 
of  the  cyclic  series. 

Crude  petroleum  is  often  used  as  a  fuel.  Its  calorific  value  is 
about  one  half  greater  than  that  of  the  best  quality  of  coal. 
Petroleum  is  refined  chiefly  by  fractional-  distillation.  It  is  also 
treated  with  concentrated  sulfuric  acid  to  remove  compounds 
with  a  disagreeable  odor  or  objectionable  properties.  Com- 
pounds of  sulfur  are  removed  by  boiling  it  with  copper  oxide. 
The  principal  product  is  usually  kerosene,  used  as  a  burning 
oil  in  lamps.  The  low-boiling  products  are  called  petroleum 
ether  and  ligroin  in  chemical  laboratories,  or,  commercially, 
gasoline,  benzene  and  naphtha,  partly  in  accordance  with  the  boil- 


290  A  TEXTBOOK  OF  CHEMISTRY 

ing  point,  but  chiefly  according  to  the  use  made  of  the  material. 
The  vapors  of  the  low  boiling  products  form  dangerously  explo- 
sive mixtures  with  air.  Kerosene  should  not  give  enough  vapor 
to  explode  at  any  temperature  below  65°  (150°  F.),  and  this  is  the 
legal  flashing  point  in  most  states. 

Products  boiling  at  a  higher  temperature  than  kerosene  are 
used  as  lubricating  oils.  Solid  products,  called  paraffin,  are 
made  into  candles  and  are  used  for  covering  jellies  and  for  many 
other  purposes.  A  semisolid  product,  called  vaseline,1  is  prepared 
for  medicinal  use.  All  of  the  products  are  very  complex  mix- 
tures of  hydrocarbons. 

Ethylene  or  Ethene,  C2H4.  When  ordinary  alcohol,  C2H5OH, 
and  concentrated  sulfuric  acid  are  mixed  in  such  proportion  that 
the  mixture  boils  at  140°,  ethyl  ether,  (C2H5)2O,  and  water  distill 
over  on  heating.  If  more  sulfuric  acid  is  used  (6  parts  of  sulfuric 
acid  to  1  of  alcohol  by  weight)  so  that  the  mixture  distills  or  de- 

H\  /H 

composes  at  170°  -  180°,  ethylene,     >C  =  C<      ,  is  formed.    In 

W  \H 

both  cases  we  may  consider  that  the  sulfuric  acid  removes  water 
from  the  alcohol,  but  the  mechanism  of  the  reaction  is  more  com- 
plicated than  such  a  statement  indicates. 

2  C2H5OH  -  H2O  =   C2H5— O— C2H5 

Ethyl  Ether 

C2H5OH  -  H2O  =  C2H4 

Ethylene 

Ethylene  is  a  colorless  gas,  very  slightly  lighter  than  air.  It 
has  a  sweetish  odor  and  burns  with  a  bright,  luminous  flame. 
The  difference  between  methane  and  ethylene  in  this  regard 
seems  to  depend  on  the  fact  that  methane  is  quite  stable,  even  at 
comparatively  high  temperatures,  and  does  not  readily  decom- 
pose with  the  separation  of  carbon,  while  at  the  temperature  of 
the  flame  ethylene  decomposes,  partly,  into  methane  and  carbon  : 

1  Vaseline  is  a  proprietary  name  used  by  the  Chesebrough  Manu- 
facturing Company.  The  name  used  in  the  Pharmacopoea  and 
by  other  manufacturers  is  petrolatum. 


UNSATURATED  COMPOUNDS  291 

f~  C 


The  carbon  which  is  liberated  temporarily  assumes  the  solid 
form  and,  being  raised  to  a  white  heat  by  the  flame,  makes 
it  luminous.  In  spite  of  this  instability,  ethylene  is  formed 
when  any  hydrocarbon  of  the  methane  series  is  heated  to  a 
high  temperature  or,  indeed,  when  almost  any  organic  com- 
pound is  heated.  For  this  reason  it  is  always  present  in 
illuminating  gas  prepared  by  heating  coal  or  oil  and  is  one  of 
the  most  important  constituents  of  the  gas,  because  of  the  lumi- 
nous quality  of  its  flame.  It  is  formed  to  some  extent  even 
from  methane,  although  it  is  less  stable  than  methane  and  the 
reaction  : 

2  CH4  =  C2H4  +  2  H2 

, 

is  endothermic.  It  would  almost  seem  that  this  ability  of  carbon 
to  enter  into  many  reactions  in  which  heat  is  absorbed  from  sur- 
rounding objects  is  one  of  the  most  important  characteristics 
of  the  element.  It  is  doubtless  intimately  connected  with  those 
properties  of  the  carbon  atom  which  cause  the  rate  of  many  of  its 
reactions  to  be  so  slow  (p.  281).  The  result  of  this  seems  to  be 
that  the  speed  of  a  reaction  in  one  direction  or  another  often  has 
more  effect  in  determining  the  direction  of  the  reaction  than  the 
heat  evolved  or  absorbed. 

Unsaturated  Compounds.  Ethylene  Chloride  and  Ethylene 
Bromide.  Ethylene  combines  directly  with  chlorine  to  form 
ethylene  chloride,  C2H4C12,  and  with  bromine  to  form  ethylene 
bromide,  C2H4Br2.  The  process  is  known  as  addition.  These 
compounds  may  be  called,  also,  dichloroethane  and  dibromo- 
ethane  and  are  to  be  considered  as  substitution  products  of 
ethane,  C2H6.  They  illustrate  the  tendency  of  compounds 
having  double  or  triple  unions  between  carbon  atoms  to  take 
up  other  elements  and  pass  back  into  compounds  which  are 
derivatives  of  the  hydrocarbons  of  the  methane  series.  For 
this  reason  the  hydrocarbons  of  the  ethylene  and  acetylene  series 
are  called  unsaturated,  while  the  hydrocarbons  of  the  methane 


292  A  TEXTBOOK  OF  CHEMISTRY 

series  and  their  derivatives  are  called  saturated.     This  conduct 

H    H 

has  led  some  chemists  to  prefer  the  formula  H — C — C — H  for 

H     H 

ethylene  instead  of  the  usual  formula,  H — C  =  C — H.  Which- 
ever formula  is  true,  it  is  evident  that  carbon  atoms  which  are 
spoken  of  as  doubly  united  are  not  more  firmly  held  together 
than  by  a  single  union.  The  reverse  of  this  seems  to  be  true. 
The  double  union  is  a  point  of  especial  reactivity. 

Acetylene.  When  an  electric  arc  is  formed  between  carbon 
points  in  an  atmosphere  of  hydrogen,  some  acetylene,  C2H2, 
is  formed.  Acetylene  is  an  endo thermic  compound  and  de- 
composes into  carbon  and  hydrogen  with  evolution  of  heat : 

C2H2  =  2C  +  H2  +  53,000  small  calories 

for  26  grams  of  acetylene.  The  formation  of  acetylene  is  often 
used  as  an  illustration  of  the  fact  that  a  high  temperature  is 
favorable  to  the  formation  of  endothermic  compounds. 

Acetylene  is  formed  by  the  incomplete  combustion  of  ethylene 
or  of  carbon  compounds  generally  and  so  is  found  among  the 
gases  coming  from  a  Bunsen  burner  burning  at  the  base.  The 
unpleasant  odor  of  these  gases  is  not,  however,  due  to  the  acety- 
lene. The  presence  of  the  acetylene  can  be  shown  by  driving 
the  gases  through  an  ammoniacal  solution  of  cuprous  chloride, 
with  which  the  acetylene  gives  a  precipitate  of  copper  carbide, 
often  incorrectly  called  copper  acetylide : 

Cu2Cl2  +  2  NH3  +  C2H2  =  Cu2C2  +  2  NH4C1 

Copper 
Carbide 

The  formation  of  copper  carbide  in  this  way  indicates  that 
acetylene  has  some  of  the  properties  of  an  acid.  The  electri- 
cal conductivity  of  solutions  of  acetylene  in  water  also  indicates 
that  it  is  an  acid,  but  an  extremely  weak  one,  so  that  its  salts 


ACETYLENE  293 

are  hydrolyzed  by  water.     On  this  fact  depends  its  prepara- 
tion for  commercial  uses  from  calcium  carbide  and  water : 

CaC2  +  2  HOH  =  Ca(OH)2  +  C2H2 

Calcium 
Carbide 

Acetylene  is  a  colorless  and  odorless  gas,  which  may  be  con- 
densed to  a  liquid  or  solid  by  cold  or  pressure.  The  boiling 
point  is  —  83.6°  and  the  melting  point  a  little  higher,  —  81.5°. 
Under  atmospheric  pressure,  therefore,  it  sublimes  without 
melting. 

Acetylene  is  unsaturated  in  the  same  sense  as  ethylene  and 
may  combine  with  four  atoms  of  bromine  to  form  acetylene 
tetrabromide  or  tetrabromoethane,  C2H2Br4. 

Acetylene  dissolves  readily  in  acetone,  one  volume  of  the 
liquid  dissolving  25  volumes  of  the  gas  at  15°  under  atmos- 
pheric pressure  and  300  volumes  under  a  pressure  of  12  atmos- 
pheres. As  a  solution  containing  100  volumes  of  the  gas  for 
one  of  acetone  is  not  explosive  (see  below)  such  a  solution  can 
be  used  to  advantage  for  illuminating  purposes. 

Acetylene  burns  from  an  ordinary  jet  with  a  smoky  flame, 
due  to  the  separation  of  carbon.  When  burned  from  a  suitable 
burner  it  gives  a  very  brilliant  white  light,  developing  from 
twelve  to  fifteen  times  as  much  light  as  can  be  obtained  with 
the  same  volume  of  good  illuminating  gas  burned  with  an  ordi- 
nary burner.  The  intense  light  is  due  partly  to  the  ease  with 
which  acetylene  decomposes  into  carbon  and  hydrogen,  but 
doubtless  also  to  the  heat  developed  by  the  decomposition, 
which  aids  in  raising  the  temperature  of  the  particles  of  carbon 
to  a  white  heat.  It  is  estimated  that  while  only  about  2  per 
cent  of  the  energy  of  illuminating  gas  is  actually  effective  as 
light,  about  10  per  cent  of  the  energy  of  burning  acetylene  may 
appear  as  light. 

At  20°  acetylene  may  be  condensed  to  a  liquid  under  a  pres- 
sure of  41  atmospheres.  This  is  considerably  less  than  the 
vapor  pressure  of  liquid  carbon  dioxide  at  the  same  tempera- 


294  A  TEXTBOOK  OF  CHEMISTRY 

ture,  and  when  a  process  had  been  invented  by  means  of  which 
the  manufacture  of  calcium  carbide  in  an  electric  furnace  be- 
came commercially  possible  it  was  thought  that  liquid  acetylene 
could  be  very  conveniently  used  in  the  liquid  form,  condensed  in 
strong  steel  cylinders.  Soon  after  the  first  attempts  in  this  direc- 
tion were  made,  however,  some  unexpected  and,  at  first,  unac- 
countable explosions  occurred.  An  investigation  of  the  matter 
soon  showed  that  liquid  acetylene,  or,  indeed,  gaseous  acetylene, 
even  under  a  pressure  of  somewhat  less  than  two  atmospheres, 
may  be  exploded  by  a  glowing  wire  or  by  a  fulminating  cap. 
The  explosion  is  due,  of  course,  to  the  fact  that  acetylene  de- 
composes into  carbon  and  hydrogen  with  considerable  evolution 

°fheat:  C2H2  =  2C  +  H2 

Although  the  volume  of  the  hydrogen  is  the  same  as  that  of 
the  acetylene  the  heat  of  decomposition  increases  the  pressure 
and  is  sufficient  to  cause  the  decomposition,  when  once  started, 
to  proceed  explosively  from  one  part  of  the  liquid  or  compressed 
gas  to  another.  It  has  been  pointed  out  above  that  solutions 
of  acetylene  in  acetone  are  not  explosive,  if  the  concentration 
is  not  carried  too  far.  Such  solutions  are  now  extensively  used, 
especially  for  automobile  and  motorcycle  lights. 

At  slightly  elevated  temperatures  acetylene  polymerizes 
easily,  that  is,  it  .combines  with  itself  to  form  more  complex 
compounds,  some  of  which  are  liquid  or  solid  at  ordinary  tem- 
peratures. When  acetylene  is  generated  by  dropping  water 
on  calcium  carbide  the  heat  evolved  by  the  reaction  may  cause 
a  very  considerable  loss  by  polymerization.  For  this  reason 
those  forms  of  generators  in  which  the  carbide  is  dropped  into 
the  water  are  most  suitable. 

Benzene,  C6H6.  In  the  manufacture  of  illuminating  gas  by 
the  distillation  of  bituminous  coal  and  also  in  the  manufacture 
of  coke  by  the  methods  in  which  retorts  are  used  and  the  by- 
products are  saved,  large  quantities  of  coal  tar  are  produced. 
This  is  an  extremely  complex  mixture  from  which  many  valuable 
products  are  obtained,  chiefly  by  fractional  distillation.  Among 


ILLUMINATING  GAS  295 

these  products  are  benzene  and  its  homologues  and  some  other 
hydrocarbons,  especially  naphthalene  and  anthracene.  These 
hydrocarbons  are  very  extensively  used  in  the  manufacture  of 
the  coal-tar  dyes  and  of  many  other  compounds  which  are  used 
in  medicine  and  in  the  arts.  Benzene  may  also  be  formed  by 
the  polymerization  of  acetylene  : 

3  C2H2  = 


It  is  a  colorless  liquid  which  melts  at  5.4°  and  boils  at  80.2°. 

Illuminating  Gas.  The  manufacture  of  illuminating  gas  by 
heating  bituminous  coal,  especially  cannel  coal,  in  earthen- 
ware or  iron  retorts,  has  been  repeatedly  referred  to.  The  gases 
obtained  in  this  manner  consist  chiefly  of  a  mixture  of  hydrogen, 
methane,  carbon  monoxide,  carbon  dioxide  and  hydrogen 
sulfide,  with  a  small  per  cent  of  the  so-called  "  heavy  hydro- 
carbons." These  last  consist  of  ethylene,  C2H4,  vapor  of 
benzene,  CeH6,  a  little  acetylene,  C2H2,  and  small  amounts  of 
many  other  gases  and  vapors.  In  burning  the  gas  only  the 
heavy  hydrocarbons  decompose  to  an  appreciable  extent  with 
separation  of  carbon.  On  this  account  these  hydrocarbons  are 
sometimes  designated  as  illuminants. 

The  composition  of  the  gas  varies  greatly  both  with  the  tem- 
perature and  the  duration  of  the  heating  of  the  coal.  A  high 
temperature  favors  the  decomposition  of  the  hydrocarbons  and 
increases  the  per  cent  of  hydrogen,  reducing  the  illuminating 
power  of  the  gas,  but  it  also  greatly  increases  the  volume  of  the 
gas  obtained  from  a  given  weight  of  coal.  The  gas  which  es- 
capes from  the  coal  during  the  first  part  of  the  heating  is  also 
much  richer  in  the  heavy  hydrocarbons  than  that  given  later. 
These  facts  are  easily  understood  from  the  instability  of  the 
hydrocarbons,  at  high  temperatures. 

The  hydrogen  sulfide  in  the  gas  must  be  removed  by  passing 
it  through  boxes  or  chambers  having  a  series  of  shelves  covered 
with  slaked  lime,  Ca(OH)2,  or  more  often  by  passing  it  through 
boxes  containing  moist  ferric  hydroxide,  Fe(OH)a: 

2  Fe(OH)3  +  3  H2S  =  2  FeS  +  S  +  3  H2O 


296  A  TEXTBOOK  OF  CHEMISTRY 

The  ferrous  sulfide  passes  back  into  a  mixture  of  ferric  hydrox- 
ide and  sulfur  on  exposure  to  the  air  and  so  may  be  used  re- 
peatedly :  2  FeS  -f  3  O  +  3  H20  =  2  Fe(OH)3  +  2  S 

The  illuminating  power  of  the  gas  is  determined  by  com- 
parison with  the  light  of  a  standard  spermaceti  candle  which 
burns  120  grains  per  hour.  For  comparison,  the  gas  is  burned 
at  the  rate  of  5  cubic  feet  per  hour.  Gas  of  good  quality  should 
give  from  18  to  22  candle  power. 

The  light  which  can  be  obtained  from  a  given  quantity  of 
gas  may  be  greatly  increased  by  burning  it  from  a  Bunsen  burner 
under  a  mantle  composed  of  oxides  of  thorium  and  cerium  — 
the  Welsbach  light,  named  after  the  inventor.  The  oxides  not 
only  furnish  the  solid  substance  heated  to  a  high  temperature, 
which  is  necessary  in  almost  all  forms  of  practical  illumination, 
but  they  also  catalyze  the  reaction  of  combustion,  localizing 
the  latter  in  immediate  contact  with  the  solid  particles  and  so 
greatly  increasing  the  temperature  to  which  these  are  raised. 
(See  p.  364.)  With  the  inverted  Welsbach  burner,  gas  may  give 
more  than  ten  times  as  much  light  as  could  be  obtained  with  an 
ordinary  flat  flame. 

Oil  Gas.  When  petroleum  is  heated  to  a  high  temperature 
the  hydrocarbons  which  it  contains  are  decomposed  with  the 
formation  of  carbon,  hydrogen,  methane  and  other  hydrocar- 
bons, some  of  which  are  gaseous  at  ordinary  temperatures.  If 
the  temperature  is  high  enough  and  especially  with  the  aid  of 
certain  catalyzers,  the  final  products  are  hydrogen  and  carbon. 
The  process  has  been  proposed  as  a  method  to  obtain  hydrogen 
for  filling  balloons.  At  lower  temperatures  it  is  possible  to 
obtain  a  gas  very  rich  in  the  heavy  hydrocarbons,  and  a  gas, 
called  Pintsch-gas,  is  manufactured  in  this  way  and  compressed 
in  steel  cylinders  for  use  in  lighting  railway  coaches  and  for 
similar  purposes. 

Water  Gas.  When  steam  is  passed  over  incandescent  coal 
or  coke  a  mixture  of  carbon  monoxide,  CO,  and  hydrogen,  H2, 
is  formed :  C  +  H2O  =  CO  +  H2 


PRODUCER  GAS  297 

The  heat  of  combustion  of  the  carbon  monoxide  and  hydrogen 
is  very  much  greater  than  the  heat  of  combustion  of  the  carbon 
or  coke.  In  other  words  the  reaction  is  endothermic  in  a  very 
high  degree.  The  mass  of  incandescent  coke  cools  very  rapidly 
as  the  reaction  proceeds.  Practically,  the  reaction  is  carried 
out  intermittently.  The  coke,  contained  in  a  large  chamber, 
is  brought  to  a  white  heat  by  burning  a  part  of  it  in  a  blast  of 
air,  while  the  products  of  combustion  are  allowed  to  escape  or 
are  utilized  as  a  fuel  gas.  The  blast  of  air  is  then  shut  off, 
steam  is  turned  on  and  the  mixture  of  carbon  monoxide  and 
hydrogen,  called  "  water  gas  "  is  collected  for  use.  After  a 
few  minutes  the  mass  cools  below  the  temperature  of  rapid 
reaction.  The  steam  is  then  shut  off  and  the  heating  process 
repeated.  By  this  method  a  gas  having  about  one  half  of  the 
heating  value  of  a  good  illuminating  gas  can  be  manufactured 
very  rapidly  and  cheaply.  It  is  not  suitable  for  use  as  an  il- 
luminating gas,  since  hydrogen  burns  with  a  colorless  flame 
and  carbon  monoxide  with  a  blue  flame  which  gives  very  little 
light.  It  may  be  enriched,  however,  by  the  addition  of  oil  gas 
and  in  that  form  is  used  as  illuminating  gas  in  many  cities  of 
the  United  States.  The  most  serious  objection  to  its  use  is 
the  very  poisonous  character  of  the  carbon  monoxide  which 
it  contains.  Not  only  is  this  dangerous  because  of  the  acci- 
dental escape  of  gas  from  an  open  stopcock,  but  in  the  winter 
time  gas  may  escape  from  leaks  in  pipes  and  may  travel  for 
some  distance  beneath  the  frozen  ground  till  it  finds  an  outlet 
in  a  cellar.  The  odor  characteristic  of  illuminating  gas  is  re- 
moved by  passage  through  earth,  still  further  increasing  the 
danger.  For  these  reasons  some  states  forbid  the  manufac- 
ture of  a  gas  containing  more  than  a  stated,  small  per  cent  of 
carbon  monoxide. 

Producer  Gas.  The  water-gas  process  is  a  wasteful  one  from 
the  point  of  view  of  the  per  cent  of  the  energy  of  the  coke  finally 
obtained  in  the  water  gas,  largely  because  of  the  thick  layer 
of  fuel  which  must  be  used.  Under  such  conditions  the  carbon 
is  burned  during  the  heating  stage  only  to  carbon  monoxide, 


298  A  TEXTBOOK  OF  CHEMISTRY 

CO,  the  carbon  dioxide  which  is  formed  in  the  lower  part  of  the 
chamber  being  reduced  to  carbon  monoxide  above : 

CO2  +  C  =  2CO 
The  heats  of  combustion  involved  are  as  follows : 

Amorphous  C  +  O2  =  CO2  +  97,650  small  calories 
CO  -f  O  =  CO2  +  68,200  small  callories 
Hence,  amorphous  C  -f-  O  =  CO  -f-  29,450  small  calories. 

It  is  evident  from  this  that  less  than  one  third  of  the  heat 
energy  of  coke  is  utilized  when  it  is  burned  only  to  carbon 
monoxide.  These  heat  relations,  which  are  so  unfavorable 
to  the  economy  of  the  water-gas  process  may  be  utilized  for 
the  production  of  a  low-grade  fuel  gas,  commonly  called  "  pro- 
ducer gas,"  which  may,  under  favorable  circumstances,  retain 
from  80  to  85  per  cent  of  the  original  heat  energy  of  the  coal 
or  fuel  employed  in  its  manufacture.  A  chamber  containing 
a  thick  bed  of  fuel  has  a  blast  of  moist  air  forced  through  it  in 
such  a  manner  that  a  gas  consisting  chiefly  of  carbon  monoxide 
and  nitrogen,  with  a  little  hydrogen  is  obtained.  The  heat 
energy  of  such  a  gas  can  be  much  more  perfectly  utilized  than 
that  of  solid  fuel  for  many  metallurgical  operations,  for  the  melt- 
ing of  glass,  for  use  in  specially  constructed  gas  engines  and  for 
many  other  purposes. 

Blast-furnace  Gas.  The  reduction  of  oxides  of  iron  by  hydro- 
gen has  been  spoken  of  as  a  reversible  reaction.  The  same  is 
true  when  carbon  monoxide  is  the  reducing  agent  : 

Fe2O3  +  3  CO  ^t  2Fe  -f  3  CO2 

Whichever  reaction  is  used,  the  equilibrium  of  the  reaction 
is  so  far  to  the  left  that  the  process  can  be  successful  only  in 
the  presence  of  a  very  large  excess  of  hydrogen  or  carbon  monox- 
ide. For  this  reason  the  gases  escaping  from  the  top  of  a 
blast  furnace  (p.  541)  are  of  much  the  same  nature  as  producer 
gas,  with  the  advantage  that  the  oxygen  of  the  carbon  monox- 
ide which  they  contain  comes  partly  from  the  iron  ore  and  so 


ILLUMINATING  AND  PRODUCER  GAS 


299 


the  per  cent  of  nitrogen  may  be  lower.  This  gas  has  long  been 
utilized  for  heating  the  blast,  generating  steam,  etc.  During 
recent  years  it  is  coming  into  extensive  use  in  gas  engines. 

The  following  table  illustrates  the  composition  of  the  various 
kinds  of  gas  which  have  been  mentioned  in  this  chapter : 


ENRICHED 

PRO- 

BLAST 

COAL 

OIL 

WATER 

DUCER 

FURNACE 

GAS 

GAS 

GAS 

GAS 

GAS 

Carbon  dioxide,  COo 

1.1 



3.0 

1.5 

11.4 

Carbon  monoxide,  CO  .     .     . 

7.2 



26.1 

23.5 

28.6 

Hydrogen  H2      

49.0 

[14.6]  l 

32.1 

6.0 

2.7 

Methane  CH4                     .     . 

345 

388 

19.8 

3.0 

0.2 

"Heavy  Hydrocarbons"    .     . 

5.0 

45.5 

16.6 

Nitrogen     

3.2 

1.1 

2.4 

66.0 

57.1 

Candle  Power       

17.5 

65.0 

25.0 

1  Ethane,  C2H6. 

Luminous  Flames.     It  was  formerly  supposed  that  carbon 

separates  in  a  flame  because  the  hydrogen  is  more  easily  burned 
than  the  carbon.  A  study  of  the  equilibrium  between  the  gases 
present  in  a  flame  has  shown,  however,  that  this  is  not  the  case. 
It  is  only  the  carbon  which  results  from  the  dissociation  of 
hydrocarbons  and  which  is  momentarily  heated  to  a  very  high 
temperature,  which  gives  the  luminous  quality  to  the  flame. 
In  a  gas  flame  burning  from  a  round  opening  we  may  distinguish 
clearly  three  parts : 

1.  The  interior  cool  portion  of  unburned  gas.     The  head  of  a 
match  inserted  to  this  point  enflames  slowly  or  not  at  all. 

2.  A  mantle  of  partial  combustion  and  of  dissociation  of 
hydrocarbons,  the  luminous  part  of  the  flame. 

3.  A  blue  mantle  surrounding  the  whole  and  especially  notice- 
able at  the  base,  where  complete  combustion  to  carbon  dioxide 
and  water  occurs. 

The  same  parts  may  be  seen  in  a  candle  flame,  the  gas  or  va- 
pors in  the  interior  being  formed  by  the  heat  of  the  flame  acting 
on  the  materials  of  the  candle. 


300  A  TEXTBOOK  OF  CHEMISTRY 

Bunsen  Burner.  A  luminous  flame  may  be  made  nonlumi- 
nous  or  slightly  luminous  by  either  of  two  methods.  The  flame 
may  be  diluted  by  an  indifferent  gas  till  it  is  cooled  below  the 
temperature  of  rapid  dissociation  for  the  hydrocarbons  present, 
or  oxygen  may  be  introduced  in  such  quantity  that  the  carbon 
burns  at  once  to  carbon  monoxide.  Both  effects  are  present 
in  the  Bunsen  burner.  The  introduction  of  air  at  the  base  of 
the  burner  causes  the  combustion  to  take  place  in  two  stages. 
In  the  inner  cone,  Fig.  84,  p.  303,  the  constituents  of  the  gas  are 
partially  burned,  giving  a  mixture  of  nitrogen,  hydrogen,  car- 
bon monoxide,  carbon  dioxide  and  water  vapor.  The  last  four 
substances  are  present  above  this  cone  in  accordance  with  the 
equilibrium  of  the  reversible  reaction  : 


At  a  temperature  of  about  850°-900°  the  four  compounds 
will  be  present  in  equal  amounts,  by  volume,  if  the  three  ele- 
ments are  present  in  the  proportion  given  in  the  equation. 
Stated  otherwise,  at  this  temperature  the  reducing  power  of 
carbon  monoxide  is  equal  to  that  of  hydrogen  and  the  oxidiz- 
ing power  of  water  vapor  is  the  same  as  that  of  carbon  dioxide. 
At  lower  temperatures  the  equilibrium  is  displaced  to  the  right 
because  the  carbon  monoxide  becomes  the  stronger  reducing 
agent,  while  at  higher  temperatures  the  equilibrium  is  displaced 
to  the  left.  It  is  in  accordance  with  this  equilibrium  that 
water  gas  prepared  at  a  high  temperature  contains  very  little 
carbon  dioxide.  This  is  of  considerable  importance  for  an 
illuminating  gas,  because  a  small  per  cent  of  carbon  dioxide 
greatly  reduces  the  illuminating  power.  At  the  outer  mantle 
of  the  Bunsen  flame  the  carbon  monoxide  and  hydrogen  burn 
to  carbon  dioxide  and  water.  By  the  arrangement  shown  in 
Fig.  82  it  is  possible  to  separate  the  two  zones  of  the  Bunsen 
flame  from  each  other  in  such  a  way  that  the  gases  between  the 
two  zones  may  be  drawn  out  and  analyzed.  By  this  method  it 
has  been  shown  that  the  equilibrium  agrees  with  that  of  the 
water-gas  reaction  given  above. 


ILLUMINATING  AND  PRODUCER  GAS  301 

In  carrying  out  the  experiment  it  is  well  to  make  the  tube  of 
the  Bunsen  burner,  shown  at  the  base  of  the  figure,  about 
60  cm.  long,  to  secure  thorough  mixture  of  the  gas  and  air,  and 
the  proportion  between  the  two  must  be  properly  regulated. 
If  a  spray  of  a  solution  containing  a  lithium  and  a  copper  salt 
is  introduced  into  the  gas,  the  lower  cone  will  be  colored 
red  from  the  lithium  and  the  upper  cone  green  from  /\ 
the  copper,  because  lithium  is  oxidized  by  even  the  '  * 
small  amount  of  oxygen  in  the  water,  carbon  monoxide 
and  carbon  dioxide  of  the  lower  cone,  while  the  copper 
is  not  oxidized  till  the  freer  oxygen  supply  of  the  upper 
cone  is  reached.  Smithells  and  Ingle,  J.  Chem.  Soc. 
61,  204  (1902) ;  Smithells,  Phil.  Mag.  [5]  39,  132  (1895). 

When  the  proportion  of  air  entering  a  Bunsen  burner 
at  the  base  is  increased,  the  inner  cone  grows  shorter 
until  a  point  is  reached  where  the  flame  "  snaps  back  " 
and  burns  at  the  base.  This  is  because  the  inner  cone 
is  to  be  looked  upon  as  a  stationary  explosion  wave 
where  the  velocity  of  the  current  of  gases  upwards  is 
equal  to  the  velocity  of  the  combustion  downwards. 
An  increase  in  the  proportion  of  oxygen  increases  this 
combustion  velocity  till  it  exceeds  the  velocity  of  the 
current  of  gas. 

Explosion  Waves.  When  a  mixture  of  gases  is  ex- 
ploded the  combustion  is  not,  of  course,  instantaneous, 
but  proceeds  with  a  definite  velocity  from  the  point  of  *  ' 
ignition.  This  phenomenon  has  been  carefully  studied, 
partly  by  measurements  of  the  velocity  by  stationary  flames,  as 
suggested  above  (Michelson,  Wied.  Ann.  37,  19  (1889)),  partly 
by  measuring  the  pressures  developed  (Michael  and  Le  Chatelier, 
Ann.  des  Mines  [8]  4,  379,  599  (1883)),  partly  by  photographing 
the  flame  by  its  own  light  (Dixon,  Ber.  38,  2426  (1905),  Phil. 
Trans.  184,  97  (1893);  200,  315  (1903)).  It  has  been  found 
that  there  are  two  distinct  stages  in  the  explosion.  At  first  the 
combustion  proceeds  with  only  a  moderate  velocity  —  for  a 
mixture  of  carbon  monoxide  and  oxygen  from  20  to  91  cm.  per 


302  A  TEXTBOOK  OF  CHEMISTRY 

second,  according  to  the  composition  of  the  mixture.  The  ex- 
pansion of  the  burning  mixture  by  the  heat  compresses  the  gas 
in  front  of  the  explosion  wave.  This  adiabatic1  compression 
raises  the  temperature  of  the  gas  and,  according  to  the  prin- 
ciples of  thermodynamics,  in  cases  where  the  final  volume  is 
diminished  by  the  explosion,  lowers  the  explosion  temperature. 
If  the  apparatus  in  which  the  explosion  occurs  is  large  enough, 
a  point  will  finally  be  reached  at  which  the  compression  raises 
the  temperature  of  the  mixture  in  front,  to  its  explosion  temper- 
ature. When  this  point  is  reached,  the  velocity  suddenly 
changes  and  the  flame  proceeds  with  the  velocity  of  sound. 
The  pressure  of  the  compression  wave  may  become  ten  times 
that  of  the  original  mixture  and  it  is  estimated  that  temperatures 
of  5000°-6000°  may  be  produced.  The  velocity  of  sound  at 
5000°  would  be  about  1400  meters  per  second  and  the  velocities 
actually  measured  somewhat  exceed  this.  This  result  is  of 
considerable  importance  in  connection  with  the  design  of  gas  or 
gasoline  engines,  where  the  form  and  dimensions  of  the  explosion 
cylinder  should  be  chosen  so  as  to  avoid  the  severe  shock  which 
comes  with  the  second  type  of  explosion  waves.  It  also  explains 
the  shattering  of  a  long  eudiometer  at  the  end  farthest  from  the 
point  of  ignition.  A  glass  tube  of  moderate  thickness  will  readily 
withstand  the  explosion  of  a  mixture  of  oxygen  and  hydrogen, 
if  the  dimensions  are  such  as  to  avoid  the  explosion  wave  of  the 
second  type. 

Temperature  of  Flames.  The  temperature  of  the  flame  of  a 
Bunsen  burner  varies  from  about  300°  in  the  center,  near  the 
mouth  of  the  burner,  where  combustion  has  not  begun,  to  about 
1550°  in  the  portion  between  the  inner  cone  and  the  outside  of 
the  flame.  These  temperatures  are  shown  in  detail  in  Fig.  84. 

In  the  Meker  burner,  Fig.  83,  by  widening  the  top  of  the 
burner  and  giving  it  a  considerable  number  of  fairly  heavy 
metallic  partitions,  the  inner  cone  is  divided  into  a  number  of 

1  Adiabatic  means  without  escape  of  heat.  Here  the  compres- 
sion is  so  rapid  that  the  heat  which  results  from  the  compression 
cannot  escape. 


ILLUMINATING  AND  PRODUCER  GAS 


303 


small  and  very  short  parts.  This  brings  the  high  temperature 
of  the  upper  part  of  the  Bunsen  flame  down  close  to  the  mouth 
of  the  burner,  concentrates  the  flame  and  gives  it  a  more  uni- 
form and  somewhat  higher  temperature. 

The  temperatures  given  in  the  figures  are,  of  course,  the  tem- 
peratures of  the  flame  when  no  substance  radiating  heat  is 


1540° 


1670* 


Fig.  83 


Fig.  84 


present.  A  platinum  or  porcelain  crucible  placed  in  the  flame 
will  be  at  a  much  lower  temperature.  A  20-gram  platinum 
crucible  placed  1  cm.  above  the  Meker  burner,  with  ordinary 
gas,  will  usually  have  a  temperature  of  900°-950°. 

Blowpipe.     By  means  of  a  blowpipe  (Fig.  85)  the  flame  of  a 
candle  or  of  a  Bunsen  burner  may  be  conveniently  used  for 


304 


A  TEXTBOOK  OF  CHEMISTRY 


Fig.  85 


heating  or  for  oxidizing  or  reducing  substances  supported  on 
charcoal  or  contained  in  beads  of  borax  or  of  sodium  meta- 
phosphate  (from  microcosmic  salt).  The  interior  of  the  flame, 
especially  if  it  retains  a  slightly  luminous  character,  will  have  a 
strong  reducing  effect,  while  at  a  point  just  beyond  the  tip  of  the 

flame,  where  substances  are 
heated  and  at  the  same  time 
can  receive  oxygen  from  the 
air,  an  oxidizing  effect  will 
be  produced.  These  effects 
can  be  readily  shown  with 
litharge,  PbO,  and  metallic  lead.  Similar  effects  can  be  ob- 
tained in  the  Bunsen  flame.  The  tip  of  the  inner  cone  is 
reducing,  while  the  outer  edge  of  the  flame  is  oxidizing. 

Reversed  Flames.  Under  ordinary  conditions  a  flame  con- 
sists of  a  combustible  gas  surrounded  by  air  or  oxygen  with 
which  it  is  combining.  Because 
the  oxygen  of  the  air  is  available 
without  expense  further  than  the 
apparatus  necessary  to  utilize  it, 
we  have  become  accustomed  to 
speak  and  think  of  the  combust- 
ible gas  as  the  source  of  the  energy 
which  we  use.  We  can,  however, 
conceive  of  a  world  where  the  at- 
mosphere should  consist  of  meth- 
ane or  some  other  gas  or  gases 
which  we  call  combustible.  In 
such  a  universe  we  might  obtain 
energy  by  preparing  and  burning 
oxygen.  Such  a  condition  can 
be  illustrated  with  the  apparatus 
shown  in  Fig.  86.  When  the  illuminating  gas  entering  through 
one  tube  is  in  excess,  oxygen  entering  through  the  other  will,  if 
ignited,  burn  in  the  atmosphere  of  illuminating  gas,  while  the  ex- 
cess of  the  gas  will  burn  at  the  end  of  the  lamp  chimney  above. 


Gas 


Oxygett 


Fig.  86 


ILLUMINATING  AND   PRODUCER  GAS  305 


EXERCISES 

1.  In  what  proportion  must  oxygen  be  mixed  with  the  following 
gases,  or  vapors  considered  as  gases,  for  their  complete  combustion? 
Methane,  ethane,  ethylene,  acetylene,  butane,  benzene,  gasoline  if  it 
has  the  average  composition  of  heptane. 

2.  In  what  proportion  must  the  same  gases  or  vapors  be  mixed  with 
air  for  their  combustion  ? 

3.  The  heats  of  combustion  of  hydrogen,  amorphous  carbon  and  of 
some  carbon  compounds  are  : 

Hydrogen,  H2  +  £  O2  =  H2O  +  68,400  small  calories. 

Amorphous  carbon,  C  +  O2  =  CO2  +  97,650  small  calories. 

Methane,  CH4  +  2  O2  =  CO2  +  2  H2O  +  214,000  small  calories. 

Ethylene,  C2H4  +  3  O2  =  2  CO2  +  2  H2O  +  333,350  small  calories. 

Acetylene,  C2H2  +  2|  O2  =  2  CO2  +  H2O  +  313,800  small  calories. 

Benzene  (vapor),  C6H6  +  7J  O2  =  6  CO2  +  3  H2O  +  799,350  small 
calories. 

These  values  are  all  for  a  final  condition  of  liquid  water  at  18°. 

If  the  following  reactions  could  occur  at  ordinary  temperatures,  how 
much  heat  would  be  evolved  or  absorbed  by  each  ? 

C  +  2  H2  =  CH4 

C2H4  =  C  +  CH4 
C2H2  =  2  C  +  H2 
C2H4  =  2  C  +  2  H2 
3  C2H2  =  C6H6 
C  +  H20  =  CO  +  H2 
C  +  2  H2O  =  CO2  +  2  H2 


CHAPTER  XVIII 

OXIDES   AND    SULFIDES    OF   CARBON.      ASSIMILATION 
AND   RESPIRATION.     CYANIDES. 

Carbon  Dioxide.  The  sources  of  carbon  dioxide  in  the  air 
and  the  maintenance  of  a  small  constant  amount,  which  fur- 
nishes carbon  for  the  growth  of  plants,  have  been  referred  to  in 
a  previous  chapter.  For  laboratory  uses  carbon  dioxide  is  pre- 
pared by  the  decomposition  of  carbonates.  Two  properties  of 
carbonic  acid  lead  to  the  decomposition  of  carbonates  when 
treated  with  almost  any  of  the  common  acids.  It  is  an  exceed- 
ingly weak  acid,  the  ionization: 

H2CO3^±H+-f-HCO3- 

taking  place  to  only  a  very  trifling  degree,  even  in  dilute  solu- 
tions, and  it  is  also  very  unstable,  dissociating  so  readily  to 
carbon  dioxide  and  water  : 

/O— H 
0=C  ±  0=C=0  +  H— O— H 


that  the  acid  can  exist  only  in  solution,  or  possibly  at  very  low 
temperatures. 

The  carbonates  usually  employed  in  the  laboratory  are  cal- 
cium carbonate,  CaCO3,  and  sodium  bicarbonate,  NaHCO3. 

Carbon  dioxide  is  also  readily  prepared  by  heating  sodium 
bicarbonate,  NaHCO3,  magnesium  carbonate,  MgCO3,  or  cal- 
cium carbonate,  CaCO3.  The  first  dissociates  to  normal  sodium 
carbonate,  Na2CO3,  carbon  dioxide  and  water  at  a  compara- 
tively low  temperature.  Magnesium  carbonate  requires  a  some- 
what higher  temperature  and  calcium  carbonate  must  be  heated 
to  812°,  or  very  bright  redness,  before  the  dissociation  pressure 
of  the  carbon  dioxide  is  equal  to  atmospheric  pressure. 

306 


CARBON  DIOXIDE 


307 


Carbon  dioxide  is  a  colorless  gas  with  a  slightly  sour  taste 
and  odor.  It  may  be  condensed  to  a  liquid  by  pressure  or  to 
a  solid  by  cold.  The  solid  melts  at  —  56.4°  at  a  pressure  of  5.1 
atmospheres,  while  its  vapor  pressure  is  one  atmosphere  at  —  79°. 
The  gas  can  be  liquefied,  therefore,  only  under  pressures  greater 
than  five  atmospheres.  If  liquid  carbon  dioxide  is  allowed  to 
escape  from  a  cylinder  in  which  it  is  kept  under  pressure,  the 
evaporation  of  a  portion  of  the  liquid  will  cool  the  remainder 
below  its  freezing  point  and  the  solid  carbon  dioxide,  which  has 
a  temperature  of  —79°,  can  be  collected  in  a  small  sack  of 
closely  woven  cloth  placed  over  the  nozzle  of  the  cylinder. 
Mixtures  of  the  solid  with  alcohol,  ether  or  acetone  are  very 
effective  for  securing  low  temperatures.  The  mixture  with 
acetone  boils  at  —88°  and  a  temperature  of  —110°  can  be  ob- 
tained by  forcing  air  through  it. 

Isothermals  of  Carbon  Dioxide..  The  critical  temperature  of 
carbon  dioxide  is  31.5°.  In  the  accompanying  diagram,  Fig.  87, 


VOLUMES. 

Fig.  87 


308  A  TEXTBOOK  OF  CHEMISTRY 

pressures  are  represented  by  the  ordinates,  volumes  by  the 
abscissas  and  temperatures  by  isothermal  lines,  which  show  the 
relation  between  pressure  and  volume  for  a  given  quantity  of 
carbon  dioxide.  Such  a  diagram  shows  clearly  all  of  the  most 
important  relations  between  pressure,  volume  and  temperature 
for  a  typical  gas.  In  the  region  of  A  where  the  pressures  are 
moderate  and  the  temperatures  are  very  considerably  above 
the  temperature  of  liquefaction,  for  these  pressures,  the  volumes 
are  very  nearly  inversely  proportional  to  the  pressures,  accord- 
ing to  the  law  of  Boyle,  and  the  distance  between  the  isother- 
mals  either  parallel  to  the  axis  of  temperatures  or  parallel  to 
the  axis  of  pressures  is  closely  proportional  to  the  absolute  tem- 
perature in  accordance  with  the  law  of  Charles.  The  higher 
the  temperature  and  the  lower  the  pressure,  the  more  nearly 
are  these  laws  accurate.  In  the  region  B  there  are  two  phases, 
liquid  and  vapor,  and  as  the  total  volume  is  independent  of  the 
pressure,  the  isothermals  are  parallel  to  the  axis  of  volumes. 
From  this  point  the  region  where  the  gas  laws  are  valid  is 
reached  either  by  an  increase  in  temperature  or  a  decrease  in 
pressure.  The  latter  can  occur  only  when  the  liquid  phase  has 
disappeared.  In  the  region  C  the  isothermals  are  above  the 
critical  temperature  and,  while  there  is  a  flattening  of  the  curves, 
showing  some  tendency  to  liquefy,  there  is  no  part  parallel  to 
the  axis  of  volumes.  In  the  region  D  the  volume  is  less  than 
that  occupied  by  the  liquid  even  at  a  considerably  lower  tem- 
perature, but  it  cannot  be  said  that  the  substance  is  liquid, 
since  no  change  of  pressure,  while  the  temperature  is  constant, 
will  cause  the  separation  into  two  phases.  The  coefficient  of 
compressibility,  however,  approaches  that  of  liquids. 

Density  of  Carbon  Dioxide.  The  weight  of  a  gram  molecular 
volume  of  carbon  dioxide  is  about  44,  while  that  of  air  is  about 
28.9.  Hence  the  gas  is  a  little  more  than  one  half  heavier  than 
air.  When  the  gas  escapes  somewhat  rapidly  from  crevices  in 
the  earth,  as  is  sometimes  the  case  in  wells  and  caves  or  mines, 
it  may  accumulate  in  sufficient  amount  to  suffocate  men  or 
animals.  In  the  Grotto  del  Cano  in  Italy  the  gas  accumulates 


CARBONIC  ACID  309 

near  the  floor  of  the  cave  in  such  a  manner  that  a  dog  entering 
the  cave  is  suffocated,  while  a  man  with  his  head  higher  up  may 
escape  injury.  The  presence  of  a  suffocating  mixture  in  a  well 
can  be  detected  by  lowering  a  candle  into  it,  though  a  candle 
may  be  extinguished  where  respiration  for  some  time  is  still 
possible. 

The  conduct  of  carbon  dioxide  in  wells  and  caves  has  given  a 
•  popular  impression  that  the  gas  will  accumulate  near  the  floor  of 
a  poorly  ventilated  room,  but  owing  to  the  rapid  diffusion  of 
gases  and  because  the  conditions  of  breathing  and  flames  cause 
an  immediate  mixture  with  the  air,  no  appreciable  accumulation 
of  this  sort  can  occur. 

Aqueous  Solutions  of  Carbon  Dioxide,  Carbonic  Acid.  At 
ordinary  temperatures  water  dissolves  approximately  its  own 
volume  of  carbon  dioxide  (1.02  volume  at  15°,  0.88  volume  at 
20°).  In  accordance  with  Henry's  law  the  weight  of  the  gas 
absorbed  is  very  nearly  proportional  to  the  pressure  and  since, 
by  the  law  of  Boyle,  the  weight  of  a  given  volume  of  a  gas  is 
proportional  to  the  pressure,  water  will  absorb  its  own  volume 
of  the  gas  either  at  low  or  high  pressures.  By  saturating  water 
with  the  gas  under  high  pressures  a  liquid  is  obtained  from  which 
carbon  dioxide  escapes  with  effervescence  on  relieving  the 
pressure.  Some  mineral  waters  of  this  kind,  especially  Apolli- 
naris  water,  and  Congress  water  from  a  spring  in  Saratoga,  are 
found  in  nature,  and  similar  waters  are  prepared  artificially 
either  for  use  directly  or  as  the  basis  of  the  so-called  "  soda 
water."  The  carbonic  acid,  H^COs,  formed  by  the  union  of 
the  carbon  dioxide  with  the  water  imparts  to  it  a  sightly  sour 
taste.  It  also  gives  an  acid  reaction,  which  may  be  shown  by 
the  reddening  of  litmus  or  the  discharge  of  the  color  of  a  faintly 
ilkaline  solution  of  phenol  phthalein.  The  carbon  dioxide  is 
•ixpelled  by  boiling  and  the  acid  reaction  disappears. 

Carbonates  and  Bicarbonates.  Hard  Waters.  The  acid 
character  of  a  solution  of  carbon  dioxide  is  also  shown  by  the 
brmation  of  carbonates  or  bicarbonates  (acid  carbonates)  with 
>ases : 


310  A  TEXTBOOK  OF  CHEMISTRY 

2  NaOH  +  H2CO3  =  Na2CO3  +  2  HOH 

Sodium 
Carbonate 

Ca(OH)2  +  H2C03  =  CaCO3  +  2  HOH 
NaOH  +  H2CO3  =  NaHCO3  +   HOH 

Sodium 
Bicarbonate 

The  ionization  of  the  bicarbonate  ion : 

HC03-  ^±  H+  +  C03= 

is  so  very  slight  that  very  few  carbonate  ions  can  exist  in  aqueous 
solutions.  Normal  carbonates  are  therefore  hydrolyzed  by 
water,  and  their  solutions  have  an  alkaline  reaction : 

Na2CO3  ^±  2Na+  +  CO3= 
CO3=  +  HOH  ^±  HCO-T  +  OH- 
or        Na2CO3  +  HOH  ^±  Na++  Na+  +  HCO-T  +  OH~ 

Bicarbonates  may  also  be  formed  by  the  action  of  carbonic 
acid  on  carbonates : 

CaCO3  +  H2C03  =  Ca(HC03)2 

The  carbonates  of  calcium  and  magnesium  are  only  very 
slightly  soluble  in  water  (CaCO3  in  77,000  parts),  but  the  acid 
carbonates,  or  bicarbonates,  are  much  more  easily  soluble. 
Natural  waters,  which  always  contain  carbonic  acid,  partly 
absorbed  from  the  air  but  chiefly  formed  by  the  oxidation  of  the 
organic  matter  of  the  soils  under  the  influence  of  bacteria,  take 
up  calcium  carbonate  and  magnesium  carbonate  from  the  soils 
in  the  form  of  bicarbonates.  Such  waters  are  called  "  hard  " 
waters.  When  boiled,  owing  to  the  ease  with  which  the  bicar- 
bonates dissociate  into  carbon  dioxide,  water  and  the  normal 
carbonate,  such  waters  give  a  deposit  of  calcium  and  magnesium 
carbonates,  which  forms  the  scale  in  teakettles  and  steam 
boilers.  When  the  hardness  is  due  to  bicarbonate  alone,  it  is 
called  "  temporary  hardness  "  and  can  be  largely  removed  by 
boiling  the  water: 

Ca(HCO3)2  =  CaCO3  +  H2O  +  CO2 


CARBON  MONOXIDE  311 

Calcium  sulfate,  CaSO4,  which  is  also  present  in  many  natural 
waters,  is  precipitated  only  when  the  water  is  concentrated  or 
when  it  is  heated  to  a  high  temperature  under  pressure.  A  water 
containing  calcium  sulfate  is  said  to  be  "permanently  hard." 
The  precipitation  at  high  temperatures  is  due  to  the  decreased 
solubility  of  calcium  sulfate  under  these  conditions,  and  the 
scale  formed  is  particularly  adherent  and  objectionable.  Such 
waters  may  be  softened  by  the  addition  of  an  alkaline  carbonate, 
phosphate,  fluoride  or  borate,  any  one  of  which  will  precipitate 
the  calcium  as  a  nearly  insoluble  compound : 

CaSO4  +  Na2CO3  =  Na2SO4  +  CaCO3 
3  CaSO4  +  2  Na3PO4  =  Ca3(PO4)2  +  3  Na2SO4 

CaS04  +  2  NaF  =  CaF2  +  Na2SO4 
CaSO4  +  Na2B4O7  =  CaB4O7  +  Na2SO4 

Carbon  Monoxide.  The  formation  of  carbon  monoxide  in 
the  manufacture  of  water  gas  and  of  fuel  gas  has  already  been 
discussed.  The  most  familiar  occurrence  of  the  gas  is  probably 
in  the  burning  of  hard  coal,  where  the  carbon  monoxide,  formed 
by  the  reduction  of  carbon  dioxide  in  the  interior  of  the  mass 
of  coal,  burns  at  the  surface  with  a  blue  flame. 

Carbon  monoxide  is  most  easily  prepared  in  the  laboratory 
by  the  decomposition  of  oxalic  acid,  H2C2O4.  Concentrated 
sulfuric  acid  assists  in  the  decomposition,  as  a  dehydrating  agent : 

H2C2O4  =  H2O  +  CO2  +  CO 

The  carbon  dioxide  must  be  removed  from  the  mixture  by 
means  of  soda  lime  or  by  passing  the  gases  through  a  wash  bottle 
containing  of  sodium  or  potassium  hydroxide. 

Carbon  monoxide  is  a  colorless  and  odorless  gas,  which  may  be 
condensed  to  a  liquid  that  boils  at  —  190°  and  to  a  solid,  which 
melts  at  —  199°.  It  burns  with  a  characteristic  blue  flame,  and 
mixtures  of  it  with  air  or  oxygen  explode  violently.  Curiously 
enough,  however,  a  perfectly  dry  mixture  of  carbon  monoxide 
and  oxygen  will  not  explode,  and  the  gao  dried  with  phosphorus 
pentoxide  will  not  burn  in  air  or  in  oxygen  which  has  been  dried 


312  A  TEXTBOOK  OF  CHEMISTRY 

with  the  same  agent.  This  recalls  the  fact  that  dry  chlorine 
does  not  act  on  iron  or  copper.  Many  similar  facts  for  which 
there  is  at  present  no  very  satisfactory  explanation  are  known. 
The  effect  of  a  minute  trace  of  moisture  in  promoting  reactions 
suggests  some  connection  with  the  effect  of  water  in  promoting 
reactions  between  electrolytes,  but  the  two  cases  seem  to  be  very 
radically  different.  One  chemist  has  gone  so  far  as  to  say  that 
probably  no  reaction  can  occur  between  two  substances  which 
are  absolutely  pure  —  but  it  is  doubtful  if  such  a  generalization 
is  justified. 

Carbon  monoxide  is  very  poisonous.  Air  containing  one  tenth 
of  one  per  cent  is  distinctly  dangerous  if  breathed  for  any  length 
of  time,  and  a  smaller  quantity,  constantly  present,  would  un- 
doubtedly cause  chronic  poisoning.  It  seems  to  combine  with 
the  hemoglobin  of  the  blood  and  to  alter  it  in  such  a  way  that 
the  hemoglobin  is  no  longer  able  to  perform  its  proper  function 
of  combining  with  oxygen  in  the  lungs  and  giving  it  up  again  for 
the  oxidation  of  other  substances  in  the  tissues  of  the  body.  It 
acts  as  a  cumulative  poison,  and  recovery  from  its  effects  is  often 
very  slow. 

Carbon  monoxide  may  be  considered  as  the  anhydride  of 
formic  acid,  HCO2H  (or  H2CO2),  which  might  also  be  called 
carbonous  acid.  At  200°-230°,  best  under  pressure,  it  combines 
directly  with  sodium  hydroxide  to  form  sodium  formate : 

CO  +  NaOH  =  HCO2Na 

The  Cycle  of  Carbon  in  Nature.  The  carbon  dioxide  of  the 
air  furnishes  the  great  storehouse  from  which  carbon  finds  its 
way  into  living  bodies  through  the  growth  of  plants.  These  are 
able  to  use  the  energy  of  the  sunlight  for  the  reduction  of  carbon 
dioxide,  possibly  first  to  formaldehyde,  CH2O,  which  then  unites 
with  itself  to  form  starch,  (CeHuAs)^  The  nitrogen  which  is 
necessary  to  form  proteins  and  other  compounds  and  the  other 
elements  required  for  the  growth  of  plants,  especially  compounds 
of  phosphorus,  potassium,  sodium,  calcium,  magnesium,  silicon 
and  iron,  must  be  furnished  by  the  soil.  The  compounds  which 


RESPIRATION  CALORIMETER  313 

are  built  up  by  the  growth  of  plants  are  ultimately  broken  down 
and  their  carbon  is  returned  to  the  air  as  carbon  dioxide  by  one 
of  three  processes  :  1 .  Wood  or  other  vegetable  material  may  be 
burned  directly.  2.  Vegetable  substances  may  be  used  as  food 
either  by  men  or  animals.  After  digestion  and  assimilation  the 
carbon  is  sooner  or  later  oxidized  to  carbon  dioxide  and  is  re- 
turned to  the  air,  mostly  through  the  lungs.  3.  Vegetable  or 
animal  substances  exposed  to  the  action  of  bacteria  decay,  and 
the  carbon  is  converted  to  carbon  dioxide. 

Respiration  Calorimeter.  From  the  point  of  view  of  considera- 
tion of  the  transformations  of  energy,  plants  in  their  growth, 
by  the  reduction  of  carbon  dioxide  and  the  formation  of  com- 
pounds in  which  the  carbon  is  combined  with  hydrogen  and 
nitrogen  as  well  as  oxygen,  store  up  energy  received  from  the 
sun,  in  the  form  of  combustible  or  edible  carbon  compounds. 
By  burning  and  by  the  use  of  steam  engines  or  other  heat  en- 
gines this  stored  chemical  energy  is  transformed  into  heat  energy 
or  mechanical  energy  for  practical  use.  When  used  as  food  by 
men  or  animals,  the  energy  is  also  transformed  both  into  heat 
and  into  muscular  and  mechanical  energy.  The  amount  of 
energy  available  in  different  kinds  of  food  can  be  accurately 
determined  by  burning  samples  of  them  in  a  calorimeter  (p.  25). 
In  order  to  follow  the  transformations  of  energy  which  occur  in 
the  human  body,  the  respiration  calorimeter  was  devised  by 
Professor  W.  O.  Atwater  with  the  aid  of  Dr.  E.  B.  Rosa,  and 
has  been  further  developed  by  Benedict,  Langworthy,  and  others. 
Figure  88  gives  a  cross  section  of  one  form  of  calorimeter  which 
shows  the  general  construction  of  the  apparatus ;  figure  89  shows 
an  exterior  view  and  some  of  the  accessory  apparatus  for  another 
form.  The  calorimeter  consists  of  a  small  room  in  which  a  man 
may  remain  for  experiments,  which  sometimes  last  for  several 
days.  The  chamber  is  provided  with  tubes  for  the  entrance 
and  exit  of  the  air  necessary  for  respiration  and  also  with  pipes 
through  which  water  can  be  circulated  to  take  up  the  heat  gen- 
erated by  the  body.  The  amount  of  food  taken,  any  change  in 
weight  of  the  body  which  occurs,  the  amount  of  carbon  dioxide 


314 


A  TEXTBOOK  OF  CHEMISTRY 


Scale:  1  Meter 


Fig.  88.  —  Horizontal  cross-section  of  chair  calorimeter,  showing 
cross-section  of  copper  wall  at  A,  zinc  wall  at  B,  hair-felt  at  E, 
and  asbestos  outer  wall  at  F ;  also  cross-section  of  all  upright 
channels  in  the  steel  construction.  At  the  right  is  the  location 
of  the  ingoing  and  outgoing  water  and  the  thermometers.  At 
C  is  shown  the  food  aperture,  and  D  is  a  gasket  separating  the 
two  parts.  The  ingoing  and  outgoing  air  pipes  are  shown  at 
the  right  inside  the  copper  wall.  The  telephone  is  shown  at  the 
left,  and  in  the  center  of  the  drawing  is  the  chair  with  its  foot- 
rest,  G.  In  dotted  line  is  shown  the  opening  where  the  man 
enters. 


RESPIRATION  CALORIMETER 


315 


evolved  and  the  quantity  of  heat  given  out  by  the  body  are  all 
carefully  measured.  The  experiments  involve,  of  course,  many 
other  details  which  cannot  be  given  here.  The  investigations 
carried  out  with  the  calorimeter  have  demonstrated  very 


Fig.  89 

clearly  that  the  amount  of  heat  generated  when  food  is  digested 
and  oxidized  in  the  human  body  is  the  same,  within  the  limits  of 
error  of  the  experiments,  as  the  amount  of  heat  generated  by 
burning  the  same  quantity  of  food  in  a  calorimeter. 

In  some  of  the  experiments  the  man  in  the  calorimeter  ex- 
pended muscular  energy  in  driving  machinery,  using  the  pedals 
of  a  bicycle  for  the  purpose.  It  was  found  that  during  this  period 
the  amount  of  heat  given  oft7  by  his  body  was  largely  increased. 
Thus,  while  at  rest  his  body,  in  one  experiment,  gave  out  112 
large  calories  per  hour.  When  performing  mechanical  labor, 
on  the  other  hand,  339  calories  per  hour  were  evolved,  while 
the  work  done  was  12,800  kilogrammeters,  equivalent  to  49 
calories.  It  will  be  seen  from  this  that  only  about  18  per  cent 


316  A  TEXTBOOK   OF  CHEMISTRY 

of  the  additional  expenditure  of  276. calories  required  for  the 
muscular  work  was  actually  converted  into  mechanical  work. 
The  human  body,  therefore,  resembles  a  steam  engine  in  dissi- 
pating as  heat  a  large  proportion  of  the  energy  required  for  its 
operation  as  a  machine. 

It  could  not  be  shown  by  the  respiration  calorimeter  that  an 
appreciable  amount  of  energy  is  required  for  mental  work. 
In  22  experiments  with  persons  performing  mental  work,  as,  for 
instance,  in  writing  an  examination  or  making  arithmetical  cal- 
culations, only  one-half  of  one  per  cent  more  heat  was  given  out 
from  the  body  than  that  given  out  when  the  same  individual 
was  at  rest.  See  Year  Book  of  the  United  States  Department 
of  Agriculture  for  1910,  p.  307. 

The  experiments  seem  to  justify  the  conclusions  that  the 
compounds  found  in  a  living  organism  obey  exactly  the  same 
chemical  and  physical  laws  as  the  same  compounds  outside  of 
of  the  body.  It  cannot,  however,  be  said  that  the  difference 
between  living  and  dead  matter  is  simply  a  difference  in  physical 
and  chemical  properties. 

*  Carbon   Suboxide,    C3O2,    is   formed   when   malonic  acid, 
CH2(CO2H)2,  is  treated  with  phosphorus  pentoxide,  P2O5 : 

OH  HH  OH 

I      \/     I 
O=C— C— C=0-*  O=C=C=C=O  +  2H2O 

Carbon  Suboxide 

Carbon  suboxide  is  a  colorless  liquid  at  a  low  temperature.  It 
boils  at  7°  and  has  a  very  strong  tendency  to  polymerize  to  a  dark 
red  substance,  probably  a  mixture  of  compounds,  having  the 
same  composition  but  evidently  a  much  higher  molecular  weight. 

*  Carbon  Oxychloride,  or  Phosgene  (Carbonyl  Chloride)  COC12. 
When  a  mixture  of  equal  volumes  of  carbon  monoxide  and  chlo- 
rine are  exposed  to  the  action  of  sunlight,  the  two  gases  unite 
to  form  carbonyl  chloride.     It  is  often  called  phosgene  from  this 
method  of   preparation  (from  <f>u<s,  light    and    yewxw,  to   pro- 
duce).    It  is  a  colorless  gas  which  may  be  easily  condensed  to 


CARBON  BISULFIDE  317 

a  liquid  that  boils  at  8.2°.     Carbonyl  chloride  is  related  to  car- 
bonic acid  just  as  sulfuryl  chloride,  SO2C12,  is  to  sulfuric  acid. 

0-H  OH  cl 

/ 
C  =  0        -^        0=0  ; 

O—H 

Carbonic 
Acid 

As  the  chloride  of  an  acid,  it  is  hydrolyzed  by  water : 

COC12  +  2  H20  -+  2  HC1  +  H2C03  ->  CO2  +  H2O 
With  ammonia  it  gives  the  amide  of  carbonic  acid,  urea : 


v-y\~v^    y 

XC1 

^0 

OH 

Cl 

Carbonyl 
Chloride 

Sulfuric 
Acid 

Sulphuryl 
Chloride 

+  4  NH3  =  CO      +  2  NH4C1 
C1  XNH2 

Urea 

Carbon  Bisulfide.  When  sulfur  vapor  is  passed  over  heated 
charcoal,  the  two  elements  combine  to  form  carbon  bisulfide, 
CS2,  which  is  now  manufactured  in  electric  furnaces.  It  is  a 
volatile,  inflammable  liquid,  which  boils  at  46.25°.  It  takes  fire 
so  easily  that  the  vapors  ignite  when  a  glass  rod,  which  has  been 
warmed  gently,  is  held  over  a  dish  containing  the  liquid.  As  a 
mixture  of  the  vapor  with  air  is  explosive,  this  property  makes 
extreme  care  necessary  in  factories  where  the  substance  is  used. 
Carbon  bisulfide  is  a  good  solvent  for  fats  and  india  rubber,  also 
for  ordinary  phosphorus,  sulfur  and  iodine.  It  is  used  in  vul- 
canizing india  rubber,  in  the  preparation  of  rubber  cements  and 
in  extracting  grease  from  wool.  It  is  poisonous  and  is  sometimes 
used  to  kill  rats  and  ground  squirrels. 

*  Sulfocarbonates.  Carbon  bisulfide  dissolves  in  solutions 
of  alkaline  sulfides,  forming  sulfocarbonates  exactly  as  carbon 
dioxide  forms  carbonates  with  alkaline  hydroxides  : 

2  KOH  +  CO2  =  K2C03  +  H2O 

Potassium 
Carbonate 

K2S  -|-  CS2  =  K2CSs 


318  A  TEXTBOOK  OF  CHEMISTRY 

or  2  KHS  +  CS2    =     K2CS3    +    H2S 

Potassium 
Sulfocarbonate 

Potassium  sulfocarbonate  is  a  yellow  salt,  easily  soluble  in 
water.  It  is  used  to  destroy  the  phylloxera  which  sometimes 
cause  great  damage  to  grapevines. 

*  Sulfocarbonic  Acid,  H2CS3,  separates  as  an  oily  liquid  when 
hydrochloric  acid  is  added  to  a  solution  of  a  sulfocarbonate.     It 
soon  decomposes  into  carbon  disulfide  and  hydrogen  sulfide. 

K2CS3  +  2  HC1  =  2  KC1  +  H2CS3  ->  H2S  +  CS2 

The  use  of  sulfocarbonates  as  germicides  depends  on  a  similar 
decomposition  by  means  of  the  carbon  dioxide  of  the  air. 

*  Carbon  Oxysulfide,  COS,  is  a  compound  intermediate  be- 
tween carbon  dioxide,  CO2,  and  carbon  bisulfide,  CS2.     It  may 
be  formed  by  the  direct  union  of  carbon  monoxide,  CO,  and  sul- 
fur, but  it  is  best  prepared  by  the  decomposition  of  a  thiocyanate 
by  means  of  an  acid  : 

KCNS    +    HC1  =  KC1  +    HCNS 

Potassium  Thiocyanic 

Thiocyanate  Acid 


H—  N4=CS  =   NH3  +  O=C=S 
H2=fO 

or         KCNS  +  2  HC1  +H2O  =  COS  +  KC1  +  NH4C1 

Carbon  oxysulfide  is  a  colorless,  odorless  gas  which  may  be 
condensed  to  a  liquid  that  boils  at  —  46.5°,  a  boiling  point  inter- 
mediate between  those  of  carbon  dioxide  and  carbon  bisulfide, 
but  much  nearer  to  the  boiling  point  of  the  former.  It  burns  in 
air  to  carbon  dioxide  and  sulfur  dioxide.  It  is  hydrolyzed  by 
water  to  thiocarbonic  acid,  which  then  decomposes  into  carbon 
dioxide  and  hydrogen  sulfide  : 

/SH  /H 

O=C=S  -f  HOH  ->  O=C          ->  O=C=O  +  S< 

\ 


Thiocarbonic 
Acid 


OXIDES  AND  SULFIDES  OF  CARBON  319 

Cyanides.  Potassium  cyanide,  KCN,  may  be  formed  by  the 
direct  union  of  the  three  elements  at  a  high  temperature,  as  when 
nitrogen  is  passed  over  a  mixture  of  carbon  and  potassium  car- 
bonate at  a  white  heat : 

K2CO3  +  4  C  +  N2  =  2  KCN  +  3  CO 

If  organic  matter  containing  nitrogen  is  heated  with  potassium 
carbonate  and  iron,  a  somewhat  similar  reaction  occurs  at  a  com- 
paratively low  temperature,  with  the  formation  of  a  double 
cyanide  of  ferrous  iron  and  potassium,  called  potassium  fer- 
rocyanide,  K4FeC6N6  (or  Fe(CN)2.4  KCN).  Cyanides  are 
formed  in  a  similar  manner  in  the  manufacture  of  illuminating 
gas  and  pass  over  into  the  ammoniacal  gas  liquors,  from  which 
they  are  recovered  commercially. 

Hydrocyanic  Acid  or  Prussic  Acid,  HCN,  may  be  obtained  by 
distilling  a  solution  of  potassium  ferrocyanide  or  potassium  cya- 
nide with  dilute  sulfuric  acid.  It  is  a  liquid  at  low  tempera- 
tures, which  boils  at  26.5°.  Hydrocyanic  acid  and  its  salts  are 
among  the  most  violent  poisons  known  and  should  always  be 
handled  with  extreme  care.  The  acid  has  a  characteristic  odor,  but 
some  persons  seem  to  be  unable  to  perceive  it ,  and  such  individuals 
need  to  be  especially  careful  in  handling  the  acid  or  cyanides. 
Hydrocyanic  acid  is  so  weak  an  acid  that  it  is  liberated  from  its 
salts  by  the  carbon  dioxide  of  the  air,  and  the  simple  cyanides 
have  the  odor  characteristic  of  the  acid. 

A  dilute  solution  of  hydrocyanic  acid  is  sometimes  used  in 
medicine. 

Potassium  Cyanide,  KCN,  or  rather  -a,  mixture  of  potassium 
and  sodium  cyanides,  is  manufactured  commercially  by  heating 
potassium  ferrocyanide  with  metallic  sodium  : 

K4FeC6N6  +  2  Na  =  4  KCN  +  2  NaCN  +  Fe 

The  salt  is  used  for  many  purposes  in  the  laboratory,  es- 
pecially in  the  preparation  of  organic  compounds.  It  is  also 
used  for  the  extraction  of  gold  from  its  ores  and  in  silver  plating. 

Complex  Cyanides.  The  cyanides  of  the  heavy  metals,  such 
as  silver,  iron,  zinc,  etc.,  are  most  of  them  insoluble  in  water,  but 


320  A  TEXTBOOK  OF   CHEMISTRY 

many  of  them  will  dissolve  in  a  solution  of  potassium  cyanide, 
ICCN.  In  the  solutions  obtained  in  this  manner  the  atoms  of 
silver  or  iron  no  longer  conduct  themselves  as  ions.1  They  will 
give  no  precipitate  with  reagents  which  precipitate  them  from 
their  ordinary  salts,  as  sodium  chloride  for  silver  or  ammonium 
sulfide  for  iron.  When  an  electric  current  passes  through  such 
?.  solution,  the  silver  or  iron  travels  with  the  cyanide  ion  toward 
the  anode,  while  only  the  potassium  goes  toward  the  cathode. 
Finally,  if  the  solutions  are  evaporated,  definite  crystalline 
compounds,  potassium  argenticyanide,  KAgC2N2,  and  potas- 
sium ferrocyanide,  K^FeCeNe.S  H2O,  can  be  obtained.  These 
facts  indicate  that  such  solutions  contain  complex  argenticyanide 
(AgC2N2~)  and  ferrocyanide  (FeCeNe")  ions  which  hold  together 
in  solution  as  the  elements  of  the  sulfate  (SO4=)  or  nitrate  (NO3~) 
ions  do.  In  further  agreement  with  this  interpretation,  such 
solutions  give  with  various  solutions  precipitates  containing  the 
characteristic  complex  group.  A  considerable  number  of  fer- 
rocyanides,  for  instance,  may  be  obtained  in  this  way. 

Ferric  salts  give  with  potassium  cyanide,  potassium  ferricy- 
anide,  K3FeC6N6  (or  FeC3N3.3  KCN),  a  salt  which  forms  red 
crystals,  while  the  crystals  of  the  ferrocyanide  are  yellow.  The 
ferrocyanide  may  be  easily  oxidized  to  the  ferricyanide  and  the 
latter  can  be  reduced  to  the  ferrocyanide  : 

2  K4FeC6N6  +  O  +  H2O  =  2  K3FeC6N6  +  2  KOH 
K3FeC6N6  +  H  +  KOH  =  K4FeC6N6  +  H2O 

Potassium  ferrocyanide  gives  with  ferric  salts  a  blue  precipi- 
tate of  Prussian  blue : 

3  K4FenC6N6  +  4  FemCl3  =  Fe4III(FeIIC6N6)3  +  12  KC1 

Prussian  Blue 

Potassium  ferricyanide  gives  a  similar  deep  blue  precipitate 
with  ferrous  salts,  but  this  is  of  variable  composition  and  usually 
contains  potassium. 

1  A  very  small  number  of  ions  of  these  metals  are  doubtless 
present.  In  some  solutions  of  this  type  hydrogen  sulfide  will  give 
a  precipitate,  owing  to  the  extreme  insolubility  of  the  sulfide  formed. 


COMPLEX  CYANIDES  .        321 

The  ferric  ferrocyanide  is  decomposed  by  potassium  hydroxide 
with  separation  of  ferric  hydroxide  while  potassium  ferrocyanide 
passes  into  solution. 

When  a  smaller  amount  of  ferric  chloride  is  added  to  a  solution 
of  potassium  ferrocyanide  a  deep  blue  solution  containing  potas- 
sium ferric  ferrocyanide  is  obtained  : 

K4FenC6N6  +  FemCl3  =  KFemFenC6  N6  +  3  KC1 

Potassium  Ferric 
Ferrocyanide 

In  this  solution  a  ferrous  salt  gives  a  precipitate  of  Turnbull's 
blue : 

2  KFem  Fe11  C6N6  -f  Fe11  C12  =  Fe11  (Fem  Fe11  C6N6)2  +  2  KC1 

Turnbull's  Blue 

The  empirical  formula  of  Prussian  blue  is  FeyCigNis,  while 
that  of  Turnbull's  blue  is  FesC^N^.  The  empirical  formula  of 
ferrous  ferricyanide  is  also  FesCi^N^. 

Both  Prussian  blue  and  Turnbull's  blue  are  used  as  blue  pig- 
ments and  for  "  blueing  "  for  laundry  purposes. 

A  solution  of  potassium  argenticyanide,  KAgC2N2  is  used  in 
manufacturing  silver-plated  ware.  While  the  silver  is  trans- 
ferred toward  the  anode  through  the  solution,  it  is  also  deposited 
on  the  cathode,  and  the  fact  that  there  are  very  few  silver  ions 
in  the  solution  for  some  reason  causes  a  smooth,  coherent  de- 
posit, while  the  silver  deposited  from  a  solution  of  silver  nitrate 
usually  assumes  a  crystalline  form.  From 'the  silver  anode, 
which  is  used  in  the  electrolysis,  silver  passes  into  solution,  re- 
generating the  potassium  argenticyanide  which  is  decomposed 
at  the  cathode. 

*  Potassium  Cyanate,  KCNO,  may  be  prepared  by  heating 
potassium  cyanide  with  lead  oxide.     It  is  very  poisonous : 

KCN  +  PbO  =  KCNO  +  Pb 

*  Potassium  Thiocyanate,  KCNS,  is  formed  when  a  mixture 
of  potassium  cyanide  and  sulfur  is  heated.     Potassium  thio- 
cyanate  and  ammonium  thiocyanate,  NH4CNS,  are  used  in 
testing  solutions  for  the  presence  of  iron  in  the  ferric  form  because 


322       .  A  TEXTBOOK  OF  CHEMISTRY 

of  the  intense  red  color  of  solutions  containing  ferric  thiocyanate, 
Fe(CNS)3. 

Cyanogen,  C2N2.  When  mercuric  cyanide,  HgC2N2,  is  heated, 
it  decomposes  into  mercury  and  cyanogen,  C2N2,  just  as  mercuric 
oxidede  composes  into  mercury  and  oxygen.  Cyanogen  is  a 
colorless,  poisonous  gas,  which  burns  with  a  characteristic  pink 
flame. 

EXERCISES 

1.  Carbon  dioxide  melts  at  —  56.4°  under  a  pressure  of  5.1  atmos- 
pheres.    Under  a  pressure  of  515  atmospheres  (500  kg.  per  sq.  cm.)  it 
melts  at  —  47.4°.     Is  the  density  of  solid  carbon  dioxide  greater  or  less 
than  that  of  the  liquid  when  both  are  present  together  ?     (See  principle 
of  van't  Hoff-Le  Chatelier,  p.  111.) 

2.  Air  contains,  normally,  about  3  parts  in  10,000  of  carbon  dioxide. 
In  accordance  with  Henry's  law  what  weight  of  the  gas  will  be  absorbed 
by  one  liter  of  water  in  contact  with  such  air  at  20°.    (See  p.  309  for  the 
solubility  of  the  gas  in  water  at  this  temperature.) 

3.  What  volume  of  normal  sodium  hydroxide  will  be  neutralized  with 
formation  of  sodium  bicarbonate,  NaHCOs,  by  one  liter  of  water  which 
has  been  in  contact  with  ordinary  air  ? 

4.  An  effervescent  drink  is  sometimes  prepared  by  mixing  two  solu- 
tions containing  sodium  bicarbonate  and  cream  of  tartar  (HKC^Oe). 
What  is  the  equation  for  the  reaction  ? 

5.  A  water  contains  0.130  g.  per  liter  of  calcium  sulfate,  CaSO4. 
How  may  grams  of  crystallized  sodium  carbonate  (Na2CO3.10  H2O) 
per  cubic  meter  will  be  required  to  soften  the  water  ?     How  many  grams 
of  sodium  fluoride  ? 

6.  How  many  grams  of  potassium  cyanide  will  be  required  to  reduce 
15  g.  of  stannic  oxide,  SnO2,  to  metallic  tin  ? 


CHAPTER  XIX 

ALCOHOLS,   ALDEHYDES,  KETONES,   ACIDS,   FATS,  CARBO- 
HYDRATES 

THE  most  important  classes  of  compounds  of  carbon  with 
oxygen  and  hydrogen  are  given  in  the  heading  of  this  chapter. 
The  number  of  such  compounds  is  very  large  and  even  a  super- 
ficial knowledge  of  them  can  be  gained  only  by  a  study  of  their 
structure,  that  is,  by  learning  the  arrangement  of  the  atoms 
within  their  molecules.  This  is  especially  true  because  many 
different  compounds  having  the  same  composition  are  known. 
Thus  there  are  no  less  than  seventy-five  compounds  having  the 
formula  C7Hi4O2.  Compounds  like  these  which  have  the  same 
composition  but  different  properties  are  called  isomers.  The 
empirical  formula  of  such  a  compound  will  give  very  little  infor- 
mation about  its  properties,  but  a  structural  formula  often  shows 
to  a  chemist,  at  once,  many  important  relationships. 

Structural  formulas  are  established  mainly  on  the  basis  of 
three  very  simple  principles  : 

1.  Valence.     The  valence  of  the  elements.     Carbon  is  almost 
always  quadrivalent,  oxygen  is  bivalent  and  hydrogen  univalent. 
Nitrogen    may   be   trivalent   or   quinquivalent.     An   example 
of  the  use  of  the  principle  of  valence  in  explaining  and  predicting 
the  structure  of  the  hydrocarbons  has  been  given  (p.  284). 

2.  Radicals.     Groups  of  atoms  called  radicals  hold  together 
and  retain  their  order  of  arrangement  in  passing  from  one  com- 
pound to  another.     Thus  in  the  reaction  represented  by  the 
equation :  C2H6O  +  HI  =  C2H5I  +  HOH 

Ethyl  Ethyl 

Alcohol  -  Iodide 

the  group,  or  radical,  ethyl,  C2H5,  is  supposed  to  retain  the  same 
arrangement  of  its  atoms  in  ethyl  iodide  that  it  has  in  ethyl 

323 


324  A  TEXTBOOK  OF  CHEMISTRY 

alcohol  and  on  the  basis  of  this  reaction  we  usually  write  the 
formula  of  ethyl  alcohol  as  C2H5OH. 

On  the  other  hand  methyl  ether,  which  has  the  same  composi- 
tion as  ethyl  alcohol,  reacts  with  hydriodic  acid  thus  : 

C2H6O  +  HI  =  CH3I  +  CH3OH 

Methyl  Methyl        Methyl 

Ether  Iodide         Alcohol 

On  the  basis  of  this  reaction  we  give  to  methyl  ether  the 
formula  CH3  —  O  —  CH3.  Expanding  these  formulas,  using  the 
principle  of  valence,  we  have  : 

H    H  H  H 

H—  C—  C—  O—  H  and  H—  C—  O—  C—  H 


A 


Ethyl  Alcohol  Methyl  Ether 

3.  Substitution.  When  an  atom  or  group  leaves  a  compound 
and  another  atom  or  group  enters  it,  the  group  which  enters  takes 
the  place  which  was  occupied  by  the  atom  or  group  which  has  left. 
In  the  reaction  between  methyl  ether  and  hydriodic  acid  given 
above  the  group  CH3O  leaves  the  compound  and  an  iodine  atom 
takes  its  place. 

Neither  of  the  last  two  principles  is  universal  in  its  application 
and  there  is  some  uncertainty  as  to  the  valence  of  the  elements 
in  some  compounds,  but,  in  spite  of  this,  it  has  been  possible  to 
determine  the  structure  of  very  many  carbon  compounds  with 
practical  certainty.  It  is  impracticable  within  the  limited  space 
which  seems  suitable  for  this  book  to  give  the  basis  for  the  struc- 
ture assigned  to  the  compounds  mentioned,  and  in  many  cases 
only  empirical  formulas  will  be  used. 

Alcohols.  Alcohols  may  be  defined  as  hydrocarbons  in  which 
one  or  more  hydrogen  atoms  have  been  replaced  by  hydroxyl,  OH. 

Methyl  Alcohol,  CHGOH,  is  obtained  along  with  acetic  acid 
and  acetone  by  the  destructive  distillation  of  wood,  and  in  an 
impure  form  it  is  often  called  wood  spirit  or  wood  alcohol.  It  is 
now  used  chiefly  as  an  addition  to  ordinary  alcohol  to  denaturize 


ALCOHOLS  325 


it.  (See  below.)  It  is  very  much  more  poisonous  than  ordi- 
nary alcohol  and  fatalities  have,  often  occurred  from  drinking 
it  or  from  breathing  its  vapor  when  varnishes  made  with  it 
were  used  in  a  confined  space.  Blindness  sometimes  results 
from  drinking  or  breathing  it.  Methyl  alcohol  boils  at  66°  and 
has  a  specific  gravity  of  0.7931  at  15.6°. 

Ethyl  Alcohol,  C2H5OH,  is  formed  by  the  fermentation  with 
yeast  of  liquids  containing  either  ordinary  cane  sugar,  as  the 
juice  of  grapes  or  apples,  or  diluted  sirups  from  the  manufacture 
of  sugar,  or  maltose  or  glucose,  sugars  formed  by  the  action  of 
malt  on  the  starch  of  Indian  corn  or  potatoes  : 


H2O  =  4  C2H5OH  +  4  CO2 

Maltose  Alcohol 

Only  dilute  alcohol  can  be  obtained  by  fermentation  and  strong 
alcohol  is  prepared  by  fractional  distillation,  the  boiling  point  of 
alcohol  being  78.3°.  Absolute  alcohol,  or  alcohol  free  from  water, 
is  prepared  by  allowing  concentrated  alcohol  to  stand  with 
quicklime,  CaO,  which  combines  with  the  last  portions  of  the 
water.  The  action  of  the  lime  may  be  hastened  by  warming 
the  mixture.  The  specific  gravity  of  absolute  alcohol  is  0.7933 
at  15°,  referred  to  water  at  4°. 

The  most  common  beverages  containing  alcohol  are  beer 
(3.5-7.5  per  cent  of  alcohol  by  volume),  wine  (9-17  per  cent), 
cider  (3.5-7.5  per  cent),  brandy,  from  the  distillation  of  wine 
(44-55  per  cent),  whisky,  from  the  distillation  of  alcoholic 
liquids  made  from  grains  (46-55  per  cent)  and  rum  (30-50  per 
cent).  The  per  cents  by  weight  are  approximately  four  fifths  of 
these. 

Alcohol  is  extensively  used  for  burning,  as  a  solvent,  in  mak- 
ing varnishes  and  in  making  pharmaceutical  extracts,  ether  and 
other  products. 

"  Denatured  alcohol  "  is  an  alcohol  to  which  some  substance 
has  been  added  to  render  it  unsuitable  for  use  as  a  beverage. 
Wood  alcohol,  gasoline  and  bone  oil  are  the  most  common 
additions.  Such  an  alcohol  is  sold  free  of  tax  and  may  be  used 


326  A  TEXTBOOK  OF  CHEMISTRY 

for  burning,  for  the  manufacture  of  varnishes  and  for  many  other 
purposes.  The  substances  added  render  it  poisonous  and  unfit 
for  drinking  or  for  any  medicinal  use. 

Phenol  or  Carbolic  Acid,  C6H5OH,  is  obtained  from  coal  tar 
and  can  be  made  artificially  from  benzene,  C6H6.  The  pure 
compound  is  solid  at  ordinary  temperatures  but  it  liquefies  on 
the  addition  of  a  small  amount  of  water.  It  has  been  exten- 
sively used  as  a  disinfectant  and  was  the  substance  first  used 
in  antiseptic  surgery.  It  is  effective  only  when  applied  in  a 
solution.  The  boiling  point  is  high  and  the  vapor  is  never 
sufficiently  concentrated  to  be  of  service.  Pure  phenol  cor- 
rodes the  tissues  and  is  a  violent  poison.  Phenol  is  the  chief 
active  constituent  of  the  "  coal-tar  dips  "  used  in  the  care  of 
sheep. 

CH2OH 

Glycerol,1  CHOH  ,  or  C3H5(OH)3,  is  an  alcohol  containing 

CH2OH 

three  hydroxyl  groups.  It  is  obtained  as  a  by-product  in  the 
manufacture  of  soap  from  fats  (p.  332).  When  glycerol  is 
treated  with  nitric  acid  under  proper  conditions,  it  reacts  with  it, 
forming  glyceryl  nitrate  (nitroglycerin)  : 

C3H5(OH)3  +  3  HNO3  =  C3H5(NO3)3  +  3  HOH 

Glyceryl 
Nitrate 

Nitroglycerin  has  a  very  much  larger  amount  of  chemical  energy 
than  the  same  elements  combined  in  the  form  of  carbon  dioxide, 
water  and  free  nitrogen,  and  under  the  influence  of  a  denotating 
cap  of  fulminate  of  mercury  the  rearrangement  of  the  atoms  to 
the  more  stable,  gaseous  forms  takes  place  so  suddenly  as  to 
cause  a  violent  explosion.  It  is  used  chiefly  for  blasting  pur- 
poses. The  liquid  form  is  sometimes  used  directly,  but  usually 

1  Glycerol  is  commonly  known  as  "glycerin."  The  name  glycerol 
is  used  by  many  writers  and  is  to  be  preferred  because  the  ending 
"ol"  is  used  for  names  of  alcohols  in  general. 


ALDEHYDES  AND  KETONES         327 

it  is  absorbed  in  sawdust  or  some  other  porous  material  and  is 
then  known  as  dynamite. 

The  conduct  of  nitroglycerin  illustrates  in  an  unusually 
striking  manner  the  fact  that  very  many  carbon  compounds, 
perhaps  even  the  majority  of  them,  are  in  a  condition  of  unstable 
equilibrium.  The  same  elements  may  be  combined  in  some 
other  form  which  contains  less  chemical  energy.  This  seems 
to  be  because  the  atoms  do  not  separate  and  recombine  so  easily 
as  is  usual  with  the  compounds  of  other  elements.  This  slow- 
ness of  reaction  velocities  and  the  fact  that  there  are  enormous 
differences  in  the  velocity  of  different  possible  reactions,  so  that 
the  compounds  which  are  formed  seldom  represent  the  lowest 
content  of  chemical  energy,  are  certainly  very  important  factors 
in  the  formation  of  the  extraordinary  number  of  carbon  com- 
pounds. 

Aldehydes  and  Ketones.     By  the  oxidation  of  an  alcohol 

/H 

containing  the  group  — C£-H        ,  an    aldehyde  may  be    ob- 

\O— H 


tained.      For  this  the  characteristic  group    is  — C<f      ,  com- 

XH 

bined  with  a  single  carbon  atom.     In  a  similar  manner  an 

alcohol  containing  the  group    /C<f         gives  a  ketone,  which 

X)H 

has  the  group  yC=O  combined  with  two  carbon  atoms. 

Formaldehyde,     H — C^       (or    H2CO),  is  easily  prepared 

by  the  oxidation  of  methyl  alcohol.  A  solution  containing  about 
40  per  cent  of  formaldehyde  is  known,  commercially,  as  for- 
malin and  is  much  used  as  a  disinfectant.  For  disinfecting  rooms 
it  must  be  disseminated  in  the  form  of  a  spray,  or  vaporized. 
It  has  the  great  advantage  over  sulfur  dioxide,  which  was  for- 
merly used  for  the  same  purpose,  that  it  does  not  cause  the  dark- 
ening of  metallic  objects  or  injury  to  fabrics.  It  is  a  very  power- 


328  A  TEXTBOOK   OF  CHEMISTRY 

ful  germicide  and  has  been  sometimes  used  as  a  preservative 
in  milk  and  other  articles  of  food.  Such  a  use  is  very  objec- 
tionable and  is  forbidden  by  law.  Formaldehyde  is  a  powerful 
poison  and  may  produce  painful  wounds  on  the  skin. 

s° 

Benzaldehyde,  C6H5C<      ,  the  chief   constituent   of  oil  of 

\H 

bitter  almonds  and  used  as  a  flavoring  extract,  is  one  of  the 
most  common  aldehydes.  Another  aldehyde  is  Citral, 

CH3v  /CH3        .0 

\C=CH-CH2—  CH2-C=CH-C<f    , 
CH/  XH 

obtained  from  the  oil  of  lemons  and  used  as  the  starting  point 
for  the  manufacture  of  ionone,  an  artificial  substitute  for  the 
odor  and  the  flavor  of  violets. 

CH3x 
Acetone,  ^>C=O,  is  formed  along  with  methyl  alcohol 

CH3/ 

and  acetic  acid  by  the  destructive  distillation  of  wood.  It  is 
also  obtained  by  distilling  calcium  acetate  or  barium  acetate. 
It  is  used  in  the  manufacture  of  chloroform  and  as  a  solvent  for 
acetylene.  Acetone  dissolves  fats  and  is  also  used  in  the  manu- 
facture of  varnishes. 

Acids.     The  further  oxidation  of  an  aldehyde  gives  an  acid, 


for  which  the  characteristic  group  is  called  carboxyl,  —  C 

OH 

The  change  from  an  alcohol  to  an  acid  is  somewhat  similar  to 

/H 
the   change  from  hydroxyl   amine,  N(-H    ,  to   nitrous  acid, 

\OH 

.     In  both  cases  the  replacement  of  the  two  hydro- 
H 

gen  atoms  by  an  oxygen  atom  gives  acid  properties  to  the  hydro- 
gen of  the  hydroxyl  group.  Very  many  acids  are  found  among 
natural  products  and  in  articles  of  food. 


ACIDS  329 


Formic  Acid,  H  —  C<T  ,  was  first  obtained  from  ants  and 

X0—  H 

is,  in  part  at  least,  the  cause  of  the  irritation  from  the  sting  of 
the  bee.  It  is  now  manufactured  on  a  large  scale  by  absorbing 
carbon  monoxide  with  sodium  hydroxide  at  a  temperature  of 
190°-220°  : 

NaOH  +  CO  =  HCOONa 

It  is  extensively  used  in  making  dyes  and  for  other  manu- 
facturing purposes  as  a  substitute  for  acetic  acid. 

>° 
Acetic  Acid,  CH3—  Cf          (or   HC2H3O2),  is   the   acid   of 

X)H 

vinegar.  It  is  formed  by  the  action  of  Bacterium  aceti  or  "mother 
of  vinegar  "  on  alcohol,  either  slowly,  as  with  wine  or  cider 
stored  in  casks  which  are  partly  open  to  the  air,  or  more  rapidly 
by  allowing  an  alcoholic  liquid  to  trickle  over  wood  shavings 
which  have  been  inoculated  with  the  bacteria.  Good  vinegar 
contains  4  per  cent  of  acetic  acid.  Acetic  acid  is  also  formed 
by  the  destructive  distillation  of  wood,  and  the  crude  acid  ob- 
tained in  this  way  is  sometimes  called  pyroligneous  acid.  This 
crude  acid,  containing  phenol  and  other  substances  which  ree'n- 
force  the  antiseptic  properties  of  the  acetic  acid,  is  sold 
under  the  name  of  "  liquid  smoke  "  as  a  preservative  for  meat. 


c 

JNOH 

Oxalic  Acid,         Q    (or  H2C2O4),  is  found  in  sheep  sorrel  and 

C/ 
XOH 

some  other  plants,  but  for  technical  purposes  it  is  manufactured 
artifically  by  heating  sodium  formate  with  a  catalyzer  : 


CO2Na 
2  HCO2Na   =  |  +H2 

Sodium  Formate  CO2Na 
Sodium 
Oxalate 


330  A  TEXTBOOK  OF  CHEMISTRY 

Oxalic  acid  is  a  comparatively  strong  acid  and  is  quite  poisonous. 
It  is  used  in  calico  printing  and  dyeing  and  in  the  bleaching  of 
flax,  straw  and  leather. 

The  decomposition  of  oxalic  acid  into  carbon  dioxide,  carbon 
monoxide  and  water  when  heated  with  concentrated  sulfuric 
acid  has  been  given  (p.  311).  The  ammonium  salt, 
(NH4)2C2O4.H2O,  is  used  to  precipitate  calcium  because 
calcium  oxalate,  CaC2O4,  is  only  very  slightly  soluble.  An 
acid  potassium  salt  called  potassium  tetraoxalate, 
KH3(C2O4)2.2  H2O,  derived  from  the  doubled  molecule,  is 
sometimes  used  as  a  standard  for  alkalimetry. 

Lactic  Acid,  CH3CHOHCO2H  (or  HC3H6O3),  is  formed  by 
the  fermentation  of  milk  sugar  and  is  found  in  sour  milk.  It  is 
both  an  alcohol  and  an  acid,  as  is  apparent  from  its  structural 
formula. 

CHOH—  C02H 

Tartaric  Acid,  |  (or   H2C4H4O6),   is    found    in 

CHOH—  C02H 

grape  juice.  When  the  sugar  of  the  juice  ferments,  forming 
alcohol,  the  acid  potassium  salt,  KHC4H4O6,  which  is  only 
slightly  soluble  in  dilute  alcohol,  separates  in  a  crude  form  called 
argoL  From  this  the  pure  salt,  known  as  cream  of  tartar, 
and  the  free  acid  are  prepared.  Cream  of  tartar  is  used  in 
cooking  and  sometimes  to  aid  in  the  formation  of  jelly  from 
fruit  juices.  Mixed  with  sodium  bicarbonate,  NaHCO3,  and 
flour  or  starch,  it  is  used  in  making  the  more  expensive  kinds 
of  baking  powders. 

Potassium  antimonyl  tartrate,  KSbOC4H4O6,  is  called  tartar 
emetic  and  is  sometimes  used  as  an  emetic. 
CH2CO2H 

Citric  Acid,  C(OH)CO2H  (or  HaCeHsOy),  is  found  in  the  juice 


CH2CO2H 

of  lemons  and  is  the  constituent  which  gives  the  sharp,  sour 
taste  to  the  fruit. 

A    cold,     concentrated     solution     of     ammonium     citrate, 


FATS  331 

will  dissolve  some  of  the  acid  calcium  phos- 
phates which  are  insoluble  in  water,  while  it  will  not  dissolve 
tricalcium  phosphate,  Ca3(PO4)2.  Its  action  in  this  respect 
is  supposed  to  resemble  the  action  of  water  containing  carbon 
dioxide  in  the  soil.  Such  a  solution  is  used  in  the  analysis  of 
commercial  fertilizers  to  distinguish  between  phosphates  which 
are  supposed  to  be  readily  available  for  the  growth  of  crops  and 
those  which  are  only  slowly  available. 
Ammonium  Ferric  Citrate,  (NH^HFeCCeHsOy^,  is  used  in 


the  preparation  of  "  blue-print  "  paper.  The  paper  is  moistened 
with  a  solution  containing  a  mixture  of  ammonium  ferric  citrate 
and  potassium  ferricyanide.  When  dry  it  is  exposed  to  the 
action  of  light  under  a  "  negative."  The  two  salts  are  easily 
soluble  in  water,  but  in  the  light  the  ferric  salt  or  the  ferricyanide 
is  reduced  by  the  citric  acid.  In  either  case,  a  blue,  insoluble 
compound  is  formed  (p.  320).  On  washing  the  paper  with 
water  the  soluble  salts  are  removed,  while  the  portions  exposed 
to  the  light  remain  a  permanent  blue. 

Benzoic  Acid,  C6H5CO2H,  is  found  in  cranberries  and  in  some 
other  fruits.  It  is  also  manufactured  by  the  oxidation  of  toluene, 
CeHsCHa,  from  coal  tar.  Even  in  dilute  solutions  it  destroys 
or  prevents  the  growth  of  bacteria  and  it  has  been  much  used 
for  this  purpose.  At  the  present  time  there  is  some  difference 
of  opinion  among  authorities  as  to  whether  small  amounts 
of  the  acid  taken  with  food  have  an  injurious  effect  or  not.  Its 
use  as  a  substitute  for  cleanliness  or  to  disguise  inferior  products 
is,  of  course,  condemned  by  every  one. 

Palmitic,  Stearic  and  Oleic  Acids,  Fats.  The  natural  fats, 
such  as  lard,  tallow,  butter,  olive  oil,  cottonseed  oil,  linseed 
oil,  and  the  oils  found  in  nuts  and  cereals  all  contain  compounds 
in  which  the  hydrogen  of  three  molecules  of  a  monobasic  or- 
ganic acid  has  been  replaced  by  the  trivalent  radical  glyceryl, 
CsHs.  The  most  common  of  these  fatty  acids  are  palmitic 
acid,  HCi6H3iO2,  stearic  acid,  HCigHssC^,  and  oleic  acid, 
HCisHsaC^.  The  corresponding  compounds  found  in  fats 
are  palmitin,  CsH^CieHs  102)3,  stearin,  CsH^CisH^^a,  and 


332  A  TEXTBOOK  OF  CHEMISTRY 

olein,  C3H5(Ci8H33O2)3.  All  three  of  these  are  found  in  lard. 
Stearin  is  found  especially  in  tallow,  palmitin  in  palm  oil,  olein 
in  olive  oil.  The  substance  known  commercially  as  stearin 
is  in  reality  an  impure  stearic  acid.  It  is  used  in  laundries  and 
for  the  manufacture  of  candles.  Stearic  and  palmitic  acids  are 
solid  at  ordinary  temperatures.  Oleic  acid  is  a  liquid. 

Soaps.  When  fats  are  heated  with  a  concentrated  solution  of 
sodium  hydroxide,  they  are  decomposed,  forming  a  sodium  salt 
of  the  fatty  acid,  and  glycerol : 

C3H5(C18H3502)3  +  3  NaOH  =  3  NaCi8H35O2  +  C3H5(OH)3 

Stearin  Sodium 

Stearate 

This  process  is  called  saponification  and  the  sodium  salts  form 
the  chief  constituents  of  the  ordinary  soaps.  The  sodium  salts 
formed  by  the  saponification  are  separated  from  the  aqueous  solu- 
tion containing  the  glycerol  by  the  addition  of  a  concentrated 
solution  of  salt,  in  which  the  salts  of  the  fatty  acids  are  nearly 
insoluble.  From  the  aqueous  solution  the  glycerol  is  recovered 
by  evaporation  and  distillation  under  diminished  pressure. 

The  action  of  soap  depends  on  the  fact  that  water  containing 
soap  in  solution  readily  forms  an  emulsion  with  oily  or  greasy 
substances  and  so  aids  in  their  removal  from  fabrics  or  from  the 
skin. 

The  calcium  salts  of  the  fatty  acids  are  mostly  so  insoluble 
in  water  that  they  cannot  act  in  this  way;  and  when  soap  is 
used  with  a  water  containing  calcium  salts  in  solution  it  can 
have  little  effect  till  enough  of  the  soap  has  been  used  to  pre- 
cipitate all  of  the  calcium  as  calcium  stearate,  Ca(Ci8H35O2)2,  or 
in  the  form  of  similar  compounds.  The  separation  of  insoluble 
calcium  salts  when  soap  is  used  with  hard  waters  is,  of  course, 
a  familiar  experience. 

Carbohydrates.  A  large  class  of  organic  substances  found  in 
plants  is  made  up  of  compounds  which  contain  oxygen  and 
hydrogen  in  the  same  proportion  as  in  water.  On  account  of 
this  composition  these  compounds  are  called  carbohydrates, 


CARBOHYDRATES  333 

but  they  do  not  contain  oxygen  and  hydrogen  in  the  form  of 
water.  Many  of  these  compounds  contain  molecules  with  six, 
twelve,  eighteen,  or  some  multiple  of  six  carbon  atoms.  The 
most  important  of  the  carbohydrates  are  the  sugars,  dextrins, 
starch  and  cellulose. 

Cane  Sugar  or  Saccharose,  Ci2H22On,  is  found  in  the  juice 
of  the  sugar  cane,  in  sugar  beets,  in  the  sap  of  the  maple  tree 
and  in  almost  all  sweet  fruits  and  vegetables.  From  the  sugar 
cane  the  juice  is  obtained  by  pressing  the  cane  between  heavy 
steel  rollers.  The  juice  is  concentrated  under  diminished 
pressure,  to  avoid  the  decomposition  which  would  occur  if  the 
solution  were  boiled  in  the  open  air  at  atmospheric  pressure. 
The  partly  concentrated  solution  is  filtered  through  boneblack 
or  animal  charcoal  to  remove  coloring  matters,  and  is  then  evap- 
orated further  till  sufficiently  concentrated  so  that  the  sugar 
will  crystallize  on  cooling.  The  crystals  are  separated  from 
the  colored  sirup  by  means  of'  rapidly  rotating  centrifugal 
strainers.  The  manufacture  of  sugar  from  beets  differs  in 
many  important  details,  but  the  general  principles  are  the 
same.  For  maple  sugar  the  juice  is  simply  evaporated  till 
the  sirup  is  sufficiently  concentrated  to  crystallize  on  cooling, 
the  other  substances  present  being  of  such  a  character  as  to 
give  the  sugar  a  desirable  flavor.  If  purified  by  the  methods 
described  above,  maple  sugar  would  not  differ  from  the  sugar 
from  sugar  cane  or  sugar  beets. 

Cane  sugar  crystallizes  in  well-formed  monoclinic  crystals, 
seen  especially  in  rock  candy.  It  melts  at  160°,  and  if  heated 
for  a  short  time  at  that  temperature  or  a  little  higher  it  is  partly 
decomposed,  giving  a  dark  brown  substance  of  indefinite  com- 
position called  caramel. 

A  solution  of  cane  sugar  turns  the  plane  of  polarization  of  a 
ray  of  polarized  light,  which  passes  through  it,  to  the  right. 
The  degree  of  rotation  is  almost  exactly  proportional  to  the 
concentration  of  the  solution,  and  the  measurement  of  the  rota- 
tion in  specially  constructed  polarimeters  called  saccharimeters 
is  very  much  used  as  a  basis  for  the  control  of  operations  in 


334  A  TEXTBOOK  OF  CHEMISTRY 

sugar  factories  and  for  the  collection  of  duty  on  sugar  at  ports 
of  entry. 

When  warmed  for  a  short  time  with  a  dilute  acid,  cane  sugar 
is  hydrolyzed  to  a  mixture  of  glucose  and  fructose : 
Ci2H22Qn  +  H2O  =  C6Hi206  +  C6Hi2O6 

Glucose         Fructose 

Glucose  rotates  the  plane  of  polarized  light  to  the  right  but 
fructose  rotates  the  plane  in  a  greater  degree  to  the  left,  and  the 
mixture  has  a  levo-rotation.  It  is  called  for  this  reason  invert 
sugar.  The  ease  with  which  the  hydrolysis  takes  place  is  a 
source  of  very  considerable  loss  in  the  manufacture  of  sugar. 
A  similar  hydrolysis  often  occurs  in  fruit  juices  and  in  honey. 

Maltose,  Ci2H22Ou,  is  formed  along  with  maltodextrin  by  the 
action  of  the  enzyme  diastase  (p.  344)  on  starch.  Its  formation 
is  a  very  important  step  in  the  manufacture  of  alcohol  from  corn 
and  other  grains  (p.  325) .  Maltose  is  hydrolyzed  to  glucose  by 
the  action  of  dilute  acids. 

Lactose,  or  Milk  Sugar,  Ci2H22On,  is  found  in  milk  and  can 
be  obtained  from  whey  as  a  by-product  in  the  manufacture 
of  cheese.  It  forms  an  important  constituent  of  milk,  as  a  food, 
and  is  used  by  preference,  rather  than  cane  sugar,  as  an  addition 
to  cows'  milk  for  feeding  infants. 

Glucose,  CeH^Oe,  is  formed  together  with  fructose  by  the 
hydrolysis  of  cane  sugar.  It  is  also  formed  by  the  hydrolysis 
of  starch  with  dilute  acids : 

(C6HioO5)n  +  n  H2O  =  n  C6Hi2O6 

Starch  Glucose 

For  the  commercial  manufacture  dilute  sulfuric  acid  is 
usually  employed  because  on  the  subsequent  addition  of  cal- 
cium carbonate  the  acid  can  be  almost  completely  removed  as 
the  difficultly  soluble  calcium  sulfate.  Glucose,  when  pure,  has 
about  three  fifths  the  sweetening  power  of  the  same  weight  of 
cane  sugar.  It  is  used  in  the  manufacture  of  "  corn  sirup,"  in 
fruit  preserves  and  in  cheap  grades  of  candy.  The  popular 
impression  that  glucose  is  harmful  as  an  article  of  diet  seems  to 
have  no  experimental  basis.  Glucose  is  dextrorotatory  and 


CARBOHYDRATES  335 

was  formerly  called  dextrose,  a  name  still  used  by  some  au- 
thors. 

In  the  disease  called  diabetes  sugar  and  starch  of  the  food 
which  is  eaten  are  changed  to  glucose  and  eliminated  in  the 
urine  instead  of  being  assimilated  as  they  should  be.  The 
glucose  can  be  detected  by  means  of  Fehling's  solution,1  a  solu- 
tion containing  potassium  sodium  tartrate,  KNaC4H4O6,  copper 
sulfate,  CuSO4,  and  sodium  hydroxide.  Glucose  reduces  the 
copper  of  such  a  solution  to  cuprous  oxide,  Cu2O,  which  sepa- 
rates as  a  red  precipitate  when  the  mixture  is  boiled. 

Fructose,  CeH^Oe,  is  the  second  constituent  of  invert  sugar 
(see  above).  It  is  levorotatory  and  was  formerly  often  called 
levulose,  but,  since  a  second,  exactly  similar  compound,  which  is 
dextrorotatory,  is  known,  the  designation  fructose  is  preferred. 
Both  glucose  and  fructose  may  be  fermented  to  alcohol  and 
carbon  dioxide  by  the  action  of  yeast. 

Starch,  (CeHioOs)^,  is  found  in  the  form  of  granules  (Fig.  90) 
which  differ  very  markedly  in  their  organized  structure  but 
which,  so  far  as  is  known,  are  identical  in  their  chemical  com- 
position. It  is  an  important  constituent  of  potatoes,  Indian 
corn,  rice,  wheat,  tapioca  and  many  other  cereals  and  vegetables. 
When  the  flour  of  a  cereal  is  kneaded  in  a  current  of  water,  the 
fine  starch  granules  float  away,  while  most  of  the  other  con- 
stituents remain  behind.  The  residue  consists  chiefly  of  ni- 
trogenous substances  and  is  called  gluten.  From  the  water  carry- 
ing the  starch  granules  in  suspension  the  latter  will  settle  out 
on  standing  or  on  allowing  the  water  to  flow  slowly  over  tables 
very  slightly  inclined.  The  practical  manufacture  of  starch  in- 
volves, of  course,  many  other  details  which  need  not  be  given  here. 

Starch  is  the  most  important  non-nitrogenous  constituent  of 
foods.  The  granules  are  covered  with  a  thin  coating  which 

1  Fehling's  solution  may  be  prepared  by  mixing  equal  volumes  of 
two  solutions  containing:  1.  34.65  grams  of  copper  sulfate 
(CuSO4.5  H2O)  in  500  cc.  of  water;  2.  173  grams  of  Rochelle 
salt  (KNaC^Oe.HaO)  and  50  grams  of  sodium  hydroxide  in 
500  cc.  of  water.  The  copper  of  1  cc.  of  the  mixed  solution  will  be 
precipitated  by  about  0.005  gram  of  glucose. 


336 


A  TEXTBOOK  OF  CHEMISTRY 


interferes  with  their  digestion,  and  one  of  the  most  important 
effects  produced  by  baking  bread  and  cooking  cereals  and  vege- 
tables is  the  bursting  of  the  granules  by  the  combined  effect  of 
heat  and  moisture.  In  the  process  of  digestion  starch  seems  to 


Fig.  90.  —  A,  potato  starch  ; 
(x!60). 


B,  rice  starch ;   C,  wheat  starch 
After  Allen. 


be  hydrolyzed  to  glucose,  which  is  then  used  to  form  a  part  of 
larger  molecules  of  compounds  found  in  the  tissues  and  fluids  of 
the  body.  These  compounds,  in  turn,  are  evidently  easily  avail- 
able in  the  animal  economy  for  the  production  of  heat  and 
energy  by  their  oxidation  to  carbon  dioxide  and  water. 

Dextrin.  If  starch  is  moistened  with  very  dilute  nitric  acid 
and  heated  for  some  time  at  120°,  it  is  converted  into  a  soluble 
compound  or  mixture  of  compounds  called  dextrin.  Other 
forms  of  dextrin  may  be  obtained  by  the  use  of  hydrochloric 
acid,  by  heat  alone  at  a  somewhat  higher  temperature,  or  by  the 
action  of  the  diastase  of  malt.  The  dextrins  are  more  or  less 
soluble  in  water  and  are  used  for  the  preparation  of  some  kinds 


CARBOHYDRATES  337 

of  mucilage  and  for  the  backs  of  postage  stamps  and  of  gummed 
labels. 

Pectose,  Pectin.  Fruits  of  nearly  all  kinds,  especially  when 
not  fully  ripe,  contain  a  substance  called  pectose,  which  is  in- 
soluble in  water  but  which  when  boiled  with  water  in  the 
presence  of  the  fruit  acids  is  decomposed  or  hydrolyzed  with 
the  formation  of  a  soluble  compound  called  pectin.  Pectin 
forms  a  jelly  with  sugar  in  a  slightly  acid  solution.  The  best 
conditions  require  enough  sugar  to  give  a  solution  boiling  at 
about  103°  and  having  a  specific  gravity  while  hot  of  1.27  to  1.29. 
The  fruit  juice  should  contain  not  less  than  0.5  to  0.7  per  cent 
of  an  organic  acid,  calculated  as  tartaric  acid.  Usually  an 
amount  of  sugar  about  one  half  to  three  fourths  of  the  volume  of 
the  fruit  juice  is  used.  The  boiling  must  not  be  long  continued 
after  separating  the  juice  from  the  fruit,  as  this  seems  to- destroy 
the  pectin  (see  N.  E.  Goldthwaite,  J.  Ind.  and  Eng.  Chem.  1, 
333  and  2,  457 ;  also  Principles  of  Jelly  Making,  Univ.  of  111. 
Bulletin,  Vol.  9,  No.  36  (1912)). 

Pectin  may  be  precipitated  from  fruit  juices,  which  have 
been  extracted  by  cooking,  by  means  of  alcohol,  but  its  com- 
position and  its  properties  as  a  definite  compound  have  not  been 
established. 

Cellulose.  The  fiber  of  wood,  flax,  cotton,  the  outer  coatings 
of  cereals  and  many  similar  materials  consist  largely  of  an  in- 
soluble substance  having  approximately  the  same  composition 
as  starch,  (CeHioOs)^.  As  coal  and  peat  were  formed  fronj 
woody  material,  cellulose  must  be  considered  as  the  principal 
original  constituent  of  all  of  our  fuels  except  petroleum  and 
natural  gas.  In  hay  and  alfalfa  it  forms  a  very  important  con- 
stituent of  the  food  of  herbivorous  animals.  It  also  furnishes 
the  basis  for  the  manufacture  of  paper.  The  best  grades  of 
filter  paper  are  nearly  pure  cellulose. 

Paper.  The  cheapest  grades  of  paper  are  made  from  straw, 
the  better  grades  from  wood  and  the  best  from  flax,  or  from 
cotton  or  linen  rags.  Many  other  fibrous  materials  may  also 
be  used.  The  materials  are  first  treated  with  a  variety  of 


338  A  TEXTBOOK  OF  CHEMISTRY 

chemicals  to  disintegrate  them,  bleach  them  and  remove  color- 
ing matters  and  other  substances  which  are  objectionable. 
This  finally  produces  a  thin,  uniform  pulp  which  can  be  spread 
out  evenly  to  form  the  sheet  of  paper.  In  the  best  kinds 
of  paper  the  fibers  must  remain  as  long  and  as  strong  as 
possible.  It  is  also  necessary  to  remove  substances  which 
turn  brown  on  exposure  to  the  light  and  which  cause  the 
paper  to  darken  with  age. 

The  glazed  surface  of  paper,  necessary  to  prevent  the  absorp- 
tion and  spreading  of  ink,  is  obtained  by  the  application  of  rosin, 
aluminium  sulfate  and  other  substances  as  sizes. 

Gun  Cotton,  Celluloid,  Lacquers.  When  cotton  is  digested 
with  a  mixture  of  concentrated  nitric  and  sulf uric  acids  a  variety 
of  compounds  called  nitro-celluloses  are  formed,  differing  with 
the  concentration  of  the  acids  used,  the  duration  of  the  treatment 
and  the  physical  condition  of  the  material.  The  most  highly 
nitrated  product  has  the  composition  C^HnO^NOs^,  and 
is  called  hexanitrocellulose.  These  products  are  powerful  ex- 
plosives and  are  used  in  torpedoes  and  also  as  the  principal 
constituent  of  smokeless  powder. 

Less  highly  nitrated  forms  dissolve  in  a  mixture  of  alcohol 
and  ether,  forming  collodion.  The  evaporation  of  the  solvent 
leaves  the  material  as  a  thin,  elastic  film  which  is  used  to  hold 
the  silver  compound  for  the  wet  plates  in  photography.  It  is 
also  sometimes  used  to  protect  wounds  from  the  access  of 
bacteria.  Similar  products  dissolved  in  amyl  acetate  form 
excellent  lacquers.  Mixed  with  or  dissolved  in  camphor  they 
form  celluloid. 


CHAPTER  XX 

AMINES,  DYES,  ALKALOIDS,  PROTEINS,  ENZYMES,  FOODS 
AND   NUTRITION 

IF  one  or  more  of  the  hydrogen  atoms  of  ammonia  are  replaced 
by  organic  radicals,  a  great  variety  of  compounds  called  amines 
are  formed.  These  compounds  combine  directly  with  acids  to 
form  salts  and  are  often  called  organic  bases,  but  it  should  be 
remembered  that  the  true  bases  are  related  to  the  amines  in 
the  same  way  that  ammonium  hydroxide,  NKUOH,  is  related 
to  ammonia,  NH3.  The  "  strength "  of  these  bases  varies 
greatly  according  to  the  nature  of  the  radicals  which  they  con- 
tain. Thus  methyl  amine,  CH3NH2,  forms  in  aqueous  solution 
a  much  stronger  base  (CH3NH3OH)  than  ammonium  hydroxide, 
while  aniline,  CeHsNH^,  gives  a  very  much  weaker  base.  Both, 
however,  form  well-crystallized,  definite  salts,  as  methyl  am- 
monium chloride,  CH3NH3C1,  or  aniline  hydrochloride, 
C6H5NH3C1. 

Methyl  Amine,  CH3NH2.  Ammonia  combines  directly  with 
methyl  iodide,  CH3I,  to  form  methyl  ammonium  iodide, 


CH,I  +  NH, 


CH3V      ,H 

—I 

Methyl  Ammonium  Iodide 


Ciisv  B 

=     H-^NO 


When  methyl  ammonium  iodide  is  warmed  with  a  concentrated 
solution  of  sodium  hydroxide,  methyl  amine  escapes  as  a  gas, 
exactly  as  ammonia  escapes  when  ammonium  chloride  is  treated 
in  the  same  way  : 

CH3NH3I  +  NaOH  =  Nal  +  CH3NH2  +  H2O 
339 


340  A  TEXTBOOK,  OF  CHEMISTRY 

Methyl  amine  is  a  gas  with  a  disagreeable  odor  resembling 
that  of  herring  brine.  It  resembles  ammonia  in  its  general 
properties,  but  is  more  easily  combustible. 

Aniline.  By  treating  benzene,  C6H6,  with  nitric  acid  a  com- 
pound called  nitrobenzene,  C6H5NO2,  can  be  prepared.  When 
this  compound  is  reduced  by  tin  and  hydrochloric  acid,  or, 
commercially,  by  iron  and  acetic  acid,  aniline,  C6H5NH2,  is  ob- 
tained. When  pure,  aniline  is  a  colorless  oil  which  boils  at  184°. 
It  is  made  on  a  large  scale  for  use  in  the  manufacture  of  a  great 
variety  of  dyes  and  for  the  preparation  of  several  valuable  com- 
pounds used  in  medicine,  especially  of  acetanilide  (antifebrin), 

yCO CH 

CeH5NHC2H3O,  and   antipyrine   CeH5 — N<;  II 

\N(CH3)-C— CH3. 
,NHC2H3O 
Phenacetine,    CeH^  ,  may  also   be  considered   as  a 

XXiH, 
derivative  of  aniline. 

Dyes.  Till  the  middle  of  the  nineteenth  century  all  of  the 
dyes  used  for  coloring  fabrics  were  either  inorganic  compounds 
or  were  natural,  vegetable  or  animal  products.  Vegetable 
products  were  chiefly  used,  but  the  number  of  those  available 
was  comparatively  small,  the  two  of  greatest  importance  being, 
probably,  indigo  and  alizarin,  or  Turkey  red.  In  1856  Sir 
William  Perkin,  in  the  course  of  some  experiments  which  he 
tried  in  the  hope  of  obtaining  quinine  from  aniline,  discovered  a 
beautiful  violet  compound,  mauve,  which  can  be  manufactured 
by  the  oxidation  of  aniline.  During  the  next  few  years  he  estab- 
lished the  manufacture  of  mauve  on  a  commercial  basis.  This 
proved  to  be  the  starting  point  for  a  great  industry  for  the  prep- 
aration of  thousands  of  different  colors  by  synthetic  processes. 
Some  of  the  artificial  dyes  are  identical  with  those  obtained 
from  vegetable  sources.  Many  others  rival  these  in  brilliancy, 
in  being  insoluble  or  "  fast  "  when  the  fabrics  dyed  with  them 
are  washed  and  in  resisting  the  action  of  light.  Others  are  not 
so  good  as  some  of  the  natural  dyes  in  these  last  particulars. 
Colors  of  almost  every  conceivable  shade  are  now  available. 


(DYES  341 

Alizarin,  Ci4H6O2(OH)2.  Shortly  after  the  discovery  of 
mauve,  Graebe  and  Liebermann,  two  German  chemists,  showed 
that  by  distilling  alizarin,  the  coloring  matter  of  Turkey  red, 
with  zinc  dust,  it  is  reduced  to  anthracene,  a  hydrocarbon 
found  in  coal  tar : 

C14H6O2(OH)2  +  5  Zn  +  H2O  =  Ci4H10  +  5  ZnO 

Alizarin     .  Anthracene 

Soon  after  this,  methods  were  developed  for  the  manufacture 
of  alizarin  from  anthracene,  and  in  a  very  few  years  the  artificial 
product  displaced  the  natural  dye  and  the  raising  of  madder 
root,  from  which  the  dye  had  been  obtained,  practically  ceased. 

Indigo,  Ci6HioO2N2.  This  dye,  which  has  been  extensively 
used  for  many  centuries,  has  been  obtained  until  recently  almost 
exclusively  from  a  plant  growing  in  India.  In  1881  Professor 
Baeyer  in  Munich  discovered  a  method  of  making  indigo  arti- 
ficially. The  process  was  complicated,  however,  and  the  original 
material  used,  toluene,  was  too  expensive  to  allow  of  the  profit- 
able manufacture  of  the  dye.  Twenty  years  of  continuous 
study  in  University  laboratories  and  in  factories  were  required 
before  a  successful  process  was  developed.  One  process  used 
starts  with  naphthalene,  Ci0H8,  and  acetic  acid  as  the  original 
materials.  In  1901  it  was  so  far  developed  that  the  Badische 
Soda-Anilin  Fabrik  had  been  willing  to  spend  $4,500,000  in  pre- 
paring for  the  manufacture  on  a  large  scale.  Since  then  the 
amount  of  the  synthetic  indigo  produced  has  increased  each 
year,  and  it  seems  likely  that  it  will  ultimately  displace  the 
natural  product. 

Indigo  is  extremely  insoluble  in  almost  all  solvents  which  do 
not  change  it  chemically,  and  its  value  depends  very  largely  on 
this  property,  which  makes  it  a  "  fast "  color.  In  order  to  fix 
it  on  the  fiber  it  is  dissolved  by  reducing  it  with  ferrous  hydroxide 
in  the  presence  of  calcium  hydroxide  or,  of  recent  years,  with 
an  alkaline  solution  of  sodium  hyposulfite,  Na2S2O4.  The  in- 
digo white  formed  by  the  reduction  is  a  weak  acid  and  forms  a 
soluble  salt  with  the  calcium  or  sodium.  Fabrics  which  are  to 


342  A  TEXTBOOK  OF  CHEMISTRY 

be  dyed  are  dipped  in  the  alkaline  solution.  On  exposure  to  the 
air  the  indigo  white  is  oxidized  back  to  indigo,  which  remains 
firmly  attached  to  the  fiber : 

Ci6Hi002N2  +  2  Fe(OH)2  +  2  H2O 

Indigo 

=  H2C16H1002N2  +  2  Fe(OH)3 

Indigo  White 

H2Ci6Hi0O2N2  +  Ca(OH)2  =  CaCi6Hi0O2N2  +  2  H2O 

Soluble  Calcium  Salt 
of  Indigo  White 

Mordants.  Some  dyes,  as  indigo,  are  so  insoluble  that  if 
once  formed  in  contact  with  the  fiber  of  a  fabric,  they  will  not 
dissolve  and  they  produce  a  "  fast  "  color.  Other  dyes  seem  to 
combine  with  fibers  directly  to  form  insoluble  compounds. 
Such  dyes  are  called  "  substantive  "  dyes  because  they  are  in- 
dependent of  the  use  of  other  substance  required  to  fix  them 
on  the  fiber.  Substantive  dyes  which  may  be  used  for  silk  or 
wool  are  much  more  common  than  those  for  cotton  or  linen. 
Other  dyes,  which  are  called  "  adjective  "  require  a  mordant, 
with  which  they  form  an  insoluble  compound,  to  fix  them. 
The  most  common  mordants  are  aluminium  acetate,  ferric 
acetate,  potassium  dichromate  and  tannic  acid. 

Alkaloids.  There  is  a  considerable  number  of  nitrogenous 
compounds,  found  in  plants,  which  combine  with  acids  to  form 
crystalline  salts  in  the  same  way  that  the  amines  do.  Some  of 
these  are  comparatively  simple  amines,  but  most  of  them  are 
complex  in  their  structure.  Many  of  them  have  some  very 
marked  physiological  action  as  poisons  or  as  medicines.  Many 
which  are  poisonous  are  used  as  medicines  in  small  doses. 

Nicotine,  CioHi4N2,  is  a  colorless  oil  found  in  tobacco,  which 
contains  from  2  to  8  per  cent  of  the  alkaloid.  It  is  very  poison- 
ous. 

Coniine,  C8Hi7N,  the  alkaloid  of  hemlock,  is  also  a  liquid. 
It  is  historically  interesting  as  the  active  principle  of  the  fatal 
draught  taken  by  Socrates. 


ALKALOIDS.    PROTEINS  343 

Atropine,  Ci7H23O3N,  is  found  in  Atropa  belladonna.  It  is 
used  to  dilate  the  pupil  of  the  eye  and  is  an  active  poison. 

Cocaine,  Ci7H2iO4N,  is  found  in  coca  leaves.  It  is  used  to 
produce  local  anaesthesia.  A  careful  study  of  cocaine  has 
shown  that  it  is  a  derivative  of  benzoic  acid  and  the  group 
derived  from  that  acid  is  chiefly  effective  in  giving  to  it  its 
valuable  qualities.  On  the  basis  of  this  discovery  other  alka- 
loids having,  in  part,  a  similar  structure  have  been  prepared. 
Some  of  these  retain  the  anaesthetic  effect  of  cocaine  and  are 
less  poisonous. 

Morphine,  CnHigOaN.H^O,  is  the  most  important  alkaloid 
of  opium  and  is  the  chief  constituent  which  gives  to  laudanum 
and  paregoric  their  poisonous  and  sedative  qualities.  Paregoric 
also  contains  camphor  and  aromatic  oils  which  may  have  as 
much  effect  as  the  morphine.  Opium  is  obtained  from  the 
poppy. 

Quinine,  C2oH24O2N2,  is  obtained  from  Peruvian  bark.  It 
is  a  specific  in  malarial  fevers. 

Strychnine,  C2iH22O2N2,  is  found  in  Strychnos  nux  vomica.  It 
is  a  violent  poison,  producing  convulsions.  A  dose  of  0.06  gram 
is  considered  fatal.  In  small  doses  it  is  a  powerful  stimulant. 

Ptomaines.  In  the  putrefaction  of  fish  or  meat  under  the 
influence  of  bacteria  and  sometimes  in  the  putrefaction  of  vege- 
table substances,  a  variety  of  basic  compounds  called  ptomaines 
is  formed.  Some  of  these  are  poisonous,  and  illness  from 
ptomaine  poisoning  often  results  from  eating  spoiled  meats, 
especially  fish.  A  few  of  them  give  color  reactions  similar  to 
those  given  by  the  vegetable  alkaloids.  Their  presence  often 
greatly  increases  the  difficulty  of  identifying  alkaloids  in  toxical 
analysis. 

Proteins.  The  most  important  compounds  in  both  animal 
and  vegetable  organisms  seem  to  be  the  proteins.  These  con- 
tain carbon,  hydrogen,  oxygen  and  nitrogen,  usually  sulfur  and 
sometimes  phosphorus  or  iron.  The  albumen  of  the  white  of  an 
egg  is  a  nearly  pure,  typical  protein.  Proteins  form  the  larger 
part  of  muscular  fiber,  of  the  casein  of  milk  and  of  the  gluten  of 


344  A  TEXTBOOK  OF  CHEMISTRY 

flour.  The  molecular  weight  of  the  proteins  contained  in  the 
substances  just  mentioned  is  very  high  —  as  much  as  15,000,  at 
least.  When  foods  containing  proteins  are  eaten,  during  the 
process  of  digestion  they  are  hydrolyzed,  under  the  influence  of 
the  enzymes  (see  below)  of  the  digestive  fluids,  giving  less  com- 
plex proteins  called  albumoses,  and  by  further  hydrolysis 
amino  acids.  These  pass  into  solution  and  so  into  the  circula- 
tion of  the  blood  and  from  them  and  from  other  portions  of  the 
food  the  organism  reconstructs  the  proteins  which  enter  into 
the  tissues  and  fluids  of  the  body,  replacing  those  proteins 
which  are  constantly  being  oxidized  to  furnish  heat  and  energy 
for  the  organism.  In  part,  they  are  oxidized  directly  without 
being  transformed  into  tissues. 

Enzymes.  The  larger  portion  of  the  foods  which  are  eaten 
are  insoluble  in  water  and  in  a  form  which  could  not  be  directly 
assimilated.  In  the  course  of  the  digestive  tract  is  found  a 
series  of  substances,  called  enzymes,  which  act  upon  the  food 
catalytically,  hydrolyzing  it  or  changing  it  so  that  it  becomes 
soluble,  and  emulsifying  the  fats.  The  most  important  of  these 
enzymes  are  the  ptyalin  of  the  saliva,  which  changes  starch  to 
sugar,  pepsin  of  the  gastric  juice  of  the  stomach,  which,  with  the 
aid  of  hydrochloric  acid,  normally  present,  changes  the  proteins 
to  albumoses  and  renders  them  soluble,  and  trypsin  from  the 
pancreas.  The  fluids  of  the  digestive  tract  are  alternately 
alkaline  and  acid. 

Many  other  enzymes  are  known.  One  of  the  first  to  be  dis- 
covered was  the  diastase,  which  is  formed  during  the  germina- 
tion of  barley  and  which  forms  the  active  constituent  of  malt. 
It  transforms  the  starch  of  grains  into  maltose  and  dextrin  in 
the  manufacture  of  alcohol.  Yeast  cells  contain  an  enzyme, 
zymase,  which  transforms  glucose,  fructose  or  maltose  to  alcohol. 

Toxins,  Antitoxins.  In  snake  venom  and  in  a  few  plants, 
especially  in  the  castor  bean,  substances  are  found  which  seem 
to  resemble  the  enzymes,  but  which  produce  disastrous,  poison- 
ous effects  upon  the  animal  organism.  These  are  called  toxins. 
Toxins  seem  also  to  be  formed  by  the  action  of  bacteria  in  cer- 


FOODS  AND  NUTRITION  345 

tain  diseases  and  doubtless  their  formation  is  often  a  chief  factor 
in  the  progress  of  the  disease.  It  has  been  discovered  that 
animals  subjected  to  the  effect  of  such  a  toxin  develop  a  sub- 
stance which  appears  to  combine  with  it  and  render  it  harmless. 
By  use  of  this  principle  it  has  been  possible  to  prepare  antitoxins 
which  are  powerful  agents  in  combating  these  diseases. 

Urea,  CON2H4.  About  85  per  cent  of  the  nitrogen  taken  as 
food  is  eliminated  from  the  human  body  in  the  form  of  urea. 
This  may  be  considered  as  formed  by  the  union  of  carbon 
dioxide  and  ammonia,  followed  by  the  loss  of  water  : 

/NH2 

/OH          /NH2 
Cf     +  2NH3  =  cf          =CO      +H20 


NH2 

Urea  is  also  formed  by  rearrangement  when  a  solution  of 
ammonium  cyanate  is  evaporated  : 

/NH2 
NH4—  O—  C=N  -»  C^=O 

Ammonium  Cyanate  NH2 

Urea 

This  transformation  of  ammonium  cyanate  into  urea,  which 
was  discovered  by  Wb'hler  in  1828,  was  the  first  synthesis  of  an 
"  organic  "  compound  from  inorganic  materials  and  it  was  the 
beginning  of  a  long  series  of  syntheses  which  have  demonstrated 
that  many  compounds  prepared  in  the  laboratory  are  identical 
with  the  same  compounds  found  in  animals  and  vegetables. 

Nutrition.  An  adult  man  weighing  about  70  kilograms,  when 
on  an  average,  mixed  diet,  eliminates  from  his  system  16  to  18 
grams  of  nitrogen  per  day.  This  is  equivalent  to  the  consump- 
tion of  100  to  112  grams  of  digestible  protein,  which  could  be 
obtained  from  3  to  3j  liters  of  milk,  1150  to  1250  grams  of 
white  bread,  600  to  750  grams  of  fresh  fish,  or  500  to  560  grams 
of  lean  beef.  From  the  experiments  with  the  respiration  calorim- 
eter (p.  313)  it  seems  that  in  a  room  at  20°  the  body  of  an 


346  A  TEXTBOOK  OF  CHEMISTRY 

adult  weighing  70  kilograms  loses  about  2200  calories  in 
24  hours.  The  protein  referred  to  above  would  give  by  its 
oxidation  in  the  body,  only  420  to  475  calories.1  To  furnish  the 
balance  of  energy  required  by  the  body,  about  325  grams  of 
carbohydrates  (starch,  sugar,  etc.)  and  50  grams  of  fat  would  be 
required.  Theoretically  it  does  not  matter  whether  the  extra 
energy  is  supplied  by  carbohydrates  or  by  fat.  Practically, 
both  are  usually  taken  somewhat  in  the  proportions  given. 

When  engaged  in  muscular  labor  for  8  hours  a  day,  about  20 
grams  of  nitrogen  are  eliminated  and  about  1800  addition  calories 
are  given  out  as  heat  and  mechanical  energy.  This  would  re- 
quire a  total  of  about  125  grams  of  protein,  625  grams  of  carbo- 
hydrates and  100  grams  of  fat.  In  the  respiration  calorimeter, 
where  the  man  drove  a  stationary  bicycle  arranged  to  measure 
the  work  performed,  about  20  per  cent  of  the  energy  of  the 
extra  food  required  was  converted  into  mechanical  energy. 
The  remainder  was  dissipated  as  heat  given  out  from  his  body. 
This  indicates  that  the  human  body,  considered  as  a  machine, 
is  somewhat  more  efficient  than  the  best  steam  engines.  But 
the  food  required  as  the  source  of  energy  is,  of  course,  much 
more  expensive  than  coal. 

A  large  number  of  dietary  studies  have  given  results  which 
indicate  that  the  average  American  diet  for  an  adult  man 
weighing  70  kilograms  is  approximately  that  given  in  the  table 
on  opposite  page. 

The  consideration  of  the  amount  of  protein  and  of  heat  energy 
required  by  the  body,  while  undoubtedly  of  great  importance 
in  deciding  upon  the  character  of  the  diet  which  should  be 
selected,  takes  account  of  only  a  few  of  the  factors  which  ought 
to  be  considered.  About  many  of  these  factors  our  knowledge 
is  still  very  imperfect.  It  is  claimed  by  some  .writers  that  the 
human  body  may  be  maintained  in  a  state  of  health  with  the 

1  In  the  bomb  calorimeter  1  gram  of  protein  gives  about  6.55 
large  calories.  When  taken  as  a  food  a  part  of  the  protein  is  elim- 
inated in  the  form  of  uric  acid,  creatinine  and  other  compounds, 
which  may  still  be  burned  with  evolution  of  heat,  hence  the  heat  of 
oxidation  of  protein  in  the  body  is  only  about  4.2  calories  per  gram. 


FOODS  AND  NUTRITION 


347 


PROPORTIONS  OF  NUTRIENTS  FURNISHED  BY  DIFFERENT  FOOD 
MATERIALS  IN  THE  AVERAGE  AMERICAN  DIETARY 


PRO- 
TEIN 

FAT 

CARBO- 
HYDRATES 

Animal  foods  : 
Total  meats      

% 
160 

% 

29  7 

% 

588 

% 

Fish     

1  8 

3  ^ 

1  0 

Effffs 

2  1 

4  1 

2Q 

-•^feft0 
Dairy  products      
Unclassified  animal  foods    .     . 

18.4 
0.2    38.5 

10.0 
0.2 

25.7 
0.2 

3.6 
0.3 

Vegetable  foods  : 
Total  cereals    

306 

43  0 

Q  1 

61  8 

Sugar,  molasses,  etc  .... 
Legumes,  tubers,  vegetables    . 
Fruits,  including  nuts     .     .  •   . 
Unclassified  vegetable  foods    . 

5.4 
20.3 
4.4 
0.5    61.2 

8.7 
0.6 

1.0 
0.7 

17.6 
12.0 
4.3 

Miscellaneous  food  materials  . 

0.3 

0.2 

0.6 

0.4 

100.0 

100.0 

100.0 

100.0 

use  of  a  very  much  smaller  amount  of  proteins,  indeed  with 
about  60  per  cent  of  that  given  above,  and  experiments  have 
been  tried  which  tend  to  support  this  point  of  view.  It  has 
been  shown,  too,  that  the  proteins  from  different  sources  differ 
very  much  in  the  character  of  the  amino  acids  formed  by  their 
hydrolysis.  Inasmuch  as  certain  amino  acids  in  definite  quan- 
tities are  required  to  restore  the  wasted  tissues  of  the  body,  it 
seems  certain  that  some  of  the  proteins  are  much  more  suitable 
than  others  for  use  as  food.  But  inquiry  along  these  lines  is 
recent  and  has  not  proceeded  far  enough  to  lead  to  final  con- 
clusions. 

Finally,  there  are  many  of  the  inorganic  elements  which  are 
absolutely  essential  to  the  health  of  the  organism,  such  as 
sodium,  chlorine,  iron,  calcium,  silicon,  sulfur,  phosphorus  and 
even  iodine. 


CHAPTER  XXI 

SILICON,     BORON,     GERMANIUM,    TIN,    LEAD,    TITANIUM, 
ZIRCONIUM,    CERIUM,   THORIUM 

THE  atomic  weights  of  the  nonmetallic  elements,  in  round 
numbers,  arranged  in  the  order  of  the  groups  of  the  Periodic 
System,  are  as  follows.  The  most  closely  related  metallic  ele- 
ments are  also  given.  All  of  the  nonmetallic  elements  are  above 
and  to  the  right  of  the  dotted  line. 


B 

11 

C 

12 

N 

14 

0 

16 

F 

19 

He 

Ne 

4 
20 

Al 

27! 

Si 

28 

P 

31 

s 

32 

Cl 

35.5 

A 

40 

Ga 

70 

Ge 

72j 

As 

75 

Se 

79 

Br 

80 

Kr 

83 

In 

114 

Sn 

118 

Sb 

120 

JTe 

127.5 

I 

127 

Xe 

130 

Tl 

204 

Pb 

207 

Bi 

208 

— 

'1 

— 

Ni 

222 

Silicon.  Si,  28.3.  Occurrence.  Silicon  is  the  most  widely 
distributed  and  abundant  element  after  oxygen.  It  forms 
about  one  fourth  of  that  portion  of  the  earth  which  we  can 
examine.  It  is  the  characteristic  element  of  minerals  almost 
as  much  as  carbon  is  the  characteristic  element  of  living  matter 
—  though  there  are  many  minerals  which  do  not  contain  sili- 
con, and  the  number  of  silicon  compounds  is  very  much  smaller 
than  that  of  the  carbon  compounds.  „ 

Silicon  is  found  in  nature  in  the  form  of  the  dioxide,  SiO2,  in 
rock  crystal,  the  purest  form  of  quartz,  amethyst,  jasper  and 
agate,  flint,  or  chalcedony.  Silicon  dioxide  also  forms  the  prin- 
cipal constituent  of  the  sandstones  and  of  ordinary  sand.  A 
hydrated  dioxide,  containing  varying  amounts  of  water,  is  called 
opal.  Silicon  is  found  in  a  great  variety  of  silicates,  all  of  which 
may  be  considered  as  salts  of  silicic  acids,  of  which  silicon  dioxide 
is  the  common  anhydride.  Among  these  may  be  mentioned 
orthoclase,  one  of  the  feldspars,  KAlSisOs,  mica,  KH2A1 3(8104) 3, 

348 


SILICON  CARBIDE  349 

kaolin,  H2Al2(SiO4)2.H2O,  a  chief  constituent  of  clay,  asbestos, 
a  variety  of  amphibole,  Ca3Mg3(SiO4)3,  talc  or  soapstone 
(alberene),  H2Mg3Si4Oi2,  serpentine  (meerschaum),  H4Mg3Si2O9, 
garnet,  Ca3Fe2(SiO4)3,  or  Ca3Al2(SiO4)3,  topaz,  Ali2Si6O25Fio, 
tourmaline,  Al4B6Oi5.4  H2NaAl3(SiO4)3,  and  beryl,  Be3Al2(SiO3)6. 
Only  a  few  ores  of  common  metals  are  silicates,  the  most  impor- 
tant being  calamine,  Zn2SiO4. 

Preparation.  Silicon  is  never  found  free  in  nature.  It  is 
most  easily  prepared  by  heating  a  mixture  of  fine  sand  with 
magnesium : 

SiO2  +  2  Mg  =  2  MgO  +  Si 

The  silicon  obtained  in  this  way  is  an  amorphous  brown  powder 
insoluble  in  water  and  acids,  except  in  a  mixture  of  hydrofluoric 
and  nitric  acids.  It  dissolves  in  a  solution  of  sodium  hydroxide 
with  evolution  of  hydrogen  : 

Si  +  2  NaOH  +  H2O  =  Na2SiO3  +  4  H 

Sodium  Silicate 

Silicon  may  be  crystallized  from  its  solution  in  melted  zinc 
and  then  forms  brilliant  needles  having  a  metallic  luster.  It  is 
now  made  in  electric  furnaces  for  use  in  the  steel  industry. 

Hydrogen  Silicide,  SiH|.  If  fine  sand  is  heated  with  four  atoms 
of  magnesium  for  each  molecule  of  the  silicon  dioxide,  the  silicon 
combines  with  the  magnesium  to  form  magnesium  silicide, 
Mg2Si.  When  this  is  treated  with  a  dilute  acid,  hydrogen  sili- 
cide is  formed : 

Mg2Si    +   4  HC1  =  2  MgCl2  +    SiH4 

Magnesium  Hydrogen 

Silicide  Silicide 

Hydrogen  silicide  is  a  colorless  gas  which  takes  fire  spontane- 
ously on  coming  to  the  air  and  burns  to  water  and  silicon  dioxide. 
It  may  be  condensed  to  a  liquid,  which  boils  at  a  very  low  tem- 
perature. 

Silicon  Carbide.  Carborundum,  SiC.  By  heating  silicon 
dioxide  with  coke  in  an  electric  furnace  it  may  be  reduced  and 


350  A  TEXTBOOK  OF  CHEMISTRY 

the  silicon  and  carbon  unite  to  form  silicon  carbide  or  carborun- 
dum, which,  when  pure,  crystallizes  in  beautiful,  colorless 
needles : 

SiO2  +  3  C  =  SiC  +  2  CO 

Carborundum  is  the  hardest  substance  known  except  the  car- 
bide of  boron  and  the  diamond,  and  it  is  manufactured  for 
use  as  an  abrasive.  For  this  purpose  it  has  partly  displaced 
corundum,  Al2Os,  which  is  used  under  the  name  of  emery. 
Carborundum  is  also  used  as  a  refractory  material  in  the  con- 
struction of  furnaces  and  to  remove  gases  from  steel. 

Silicon  Fluoride,  SiF4.  The  formation  of  silicon  fluoride  in  the 
etching  of  glass  by  hydrofluoric  acid  has  been  referred  to.  The 
compound  is  most  easily  prepared  by  warming  a  mixture  of  sand, 
SiO2,  calcium  fluoride,  CaF2,  and  concentrated  sulfuric  acid : 

2  CaF2  +  SiO2  +  2  H2SO4  =  2  CaSO4  +  SiF4  +  2  H2O 

Silicon  fluoride  is  a  gas,  but  may  be  condensed  to  a  solid  which 
melts  at  —  97°  and  has  a  vapor  pressure  of  760  mm.  at  —  90°. 

Fluosilicic  Acid,  H2SiFe.  Silicon  fluoride  is  hydrolyzed  by 
water,  as  would  be  expected  of  a  halogen  compound  of  a  non- 
metallic  element,  but  the  hydrofluoric  acid  formed  combines 
with  some  of  the  undecomposed  silicon  fluoride  to  form  a  com- 
plex acid,  fluosilicic  acid : 

SiF4   +   4  HOH  =  Si(OH)4  +  4  HF 

Silicon  Silicic 

Fluoride  Acid 

2  HF  +  SiF4  =  H2SiF6 

Fluosilicic 
Acid 

The  potassium  salt  of  fluosilicic  acid,  K2SiF6,  is  difficultly 
soluble.  The  barium  salt,  BaSiF6,  is  also  extremely  insoluble, 
even  in  dilute  acids.  It  is  the  only  salt  of  barium  likely  to  be 
mistaken  for  barium  sulfate,  BaSO4,  when  barium  chloride, 
BaCl2,  is  used  to  test  for  sulfates  in  a  dilute  acid  solution. 


SILICON  DIOXIDE  351 


Silicon  Tetrachloride,  SiCU,  is  formed  when  chlorine  is  passed 
over  heated  silicon.  It  was  formerly  prepared  by  heating  a  mix- 
ture of  silicon  dioxide  and  charcoal  in  a  current  of  chlorine  : 

Si02  +  2  Cla  +  2  C  =  SiCl4  +  2  CO 

This  method  of  preparation,  which  was  formerly  much  used 
to  obtain  chlorides  of  elements,  such  as  silicon,  aluminium  and 
chromium,  which  cannot  be  reduced  from  their  oxides  by  carbon 
at  any  moderate  temperature,  has  ceased  to  be  of  practical 
impertance  since  other  methods  have  been  developed  for  the 
preparation  of  the  free  elements.  The  process  depends  on  the 
simultaneous  use  of  chlorine  and  carbon  to  cause  the  separation 
of  the  silicon  and  oxygen.  Silicon  tetrachloride  is  a  liquid 
which  boils  at  56.9°. 

It  is  hydrolyzed  by  .water  to  silicic  acid,  Si(OH)4,  and  hydro- 
chloric acid. 

Silicon  Hexaiodide,  Si2l6,  is  formed  when  silicon  tetraiodide, 
SiI4,  is  heated  to  290°-300°  with  powdered  silver.  It  crystallizes 
from  carbon  bisulfide  in  colorless  prisms.  It  is  hydrolyzed  by 
water  to  silicooxalic  acid  : 

Si2I6  +  4  HOH  =  H2Si2O4  +  6  HI 

Silicooxalic 
Acid 

These  compounds  and  some  others,  which  have  been  prepared, 
show  that  silicon  atoms  may  combine  together  as  carbon  atoms 
do,  but  only  a  few  such  compounds  are  known  and  most  of  these 
are  comparatively  unstable.  Silicooxalic  acid,  H2Si2O4,  is  so 
unstable  as  to  be  explosive.  It  seems  evident  that  stable  com- 
plex molecules  containing  silicon  are  formed  only  when  the  atoms 
are  held  together  through  the  agency  of  some  other  element, 
such  as  oxygen.  (See  below  under  Silicic  Acids.) 

Silicon  Dioxide  or  Silica,  SiO2.  In  addition  to  the  forms  of 
occurrence  already  given,  infusorial  earth  or  "  Kieselguhr," 
a  porous  material  composed  of  the  skeletons  of  infusoria,  may  be 
mentioned.  It  is  used  to  absorb  nitroglycerin  in  the  manu- 
facture of  dynamite  and  as  a  packing  material  for  bottles  con- 


352  A  TEXTBOOK  OF  CHEMISTRY 

taining  bromine  or  other  corrosive  chemicals.  It  is  also  used 
in  sapolio  and  in  other  scouring  soaps  and  scouring  materials. 

Clear  specimens  of  rock  crystal  are  sometimes  used  in  the 
preparation  of  prisms  and  lenses  which  are  more  transparent  than 
ordinary  glass  to  ultra-violet  light.  Flint  and  pure  sands  and 
sandstones  are  used  in  the  manufacture  of  glass. 

Silicon  dioxide  is  found  in  nature  in  two  forms.  Rock  crystal 
or  quartz,  which  is  much  the  more  common,  crystallizes  in  forms 
of  the  hexagonal  system  and  has  a  specific  gravity  of  2.6.  Quartz 
can  be  formed  only  at  temperatures  below  900°.  Above  that 
temperature  it  changes  to  tridymite,  which  crystallizes  in  the 
rhombic  system  and  has  a  specific  gravity  of  2.28. 

Quartz  melts  at  about  1750°,  but  softens  and  becomes  plastic 
at  1600°  or  below.  By  means  of  the  electric  furnace  or  the 
oxyhydrogen  flame  it  can  be  melted  and  fashioned  into  tubes, 
crucibles,  beakers,  flasks,  thermometers  and  other  laboratory 
utensils.  It  has  the  advantage  over  glass  that  it  can  be  heated 
to  very  high  temperatures  without  melting  and  also  that  it  has 
such  a  small  coefficient  of  expansion  for  a  change  of  temperature 
that  it  does  not  crack  when  subjected  to  sudden  heating  or 
cooling.  Partly  for  the  same  reason  thermometers  made  from 
it  show  no  depression  of  the  zero  point  after  use  at  high  tempera- 
tures, as  is  common  with  glass  thermometers. 

Fused  quartz  has  a  specific  gravity  of  only  2.20,  nearly  the 
same  as  that  of  tridymite.  Fused  tridymite  would  doubtless 
be  a  more  correct  name  than  fused  quartz. 

Artificial  Silicates.  When  silicon  dioxide  is  heated  with  so- 
dium carbonate  or  with  the  oxide  or  carbonate  of  almost  any 
metal,  a  silicate  is  formed : 

Na2CO3  +  SiO2  =  Na2SiO3  +  CO2 

Sodium  Sodium 

Carbonate  Silicate 

CaO  +  SiO2  =  CaSiO3 

The  industrial  importance  of  reactions  of  this  character  will 
be  seen  when  it  is  stated  that  similar  reactions  are  used  for  the 


SILICIC  ACIDS 


353 


manufacture  of  glass,  for  the  formation  of  fusible  slags  in  the 
manufacture  of  iron  and  in  other  metallurgical  operations  and 
in  the  manufacture  of  Portland  cement. 

The  silicates  of  sodium  and  potassium  are  soluble  in  water, 
and  the  former,  especially,  is  called  water  glass  and  is  used  as  an 
addition  to  laundry  soaps,  for  preserving  eggs  and  for  the  fire- 
proofing  of  wood  and  fabrics.  Almost  all  other  silicates  and 
mixed  silicates  are  insoluble  or  nearly  insoluble  in  water. 

Silicic  Acids.     If  concentrated  hydrochloric   acid  is   added 
quickly  to  a  solution  of  sodium  silicate,  the  silicate  seems  to  be 
completely  decomposed,  but  the 
silicic  acid  formed  does  not  pre- 
cipitate.    If    such    an    acidified 
solution  is  dialyzed    by  placing 
it  in  a  parchment  sack  suspended 
in  water  (Fig.   91),  the  sodium 
chloride  will  diffuse  through  the 
walls  of  the  sack,  and  by  repeat- 
edly changing  the  water  on  the 
outside  the  chloride  can  be  al- 
most completely  removed,  leaving  Fig.  91 
a  colloidal  solution  of  silicic  acid. 

Colloidal  solutions  may  also  be  prepared  by  the  hydrolysis 
of  esters  of  silicic  acid  such  as  the  methyl  ester,  (CH3)4SiO4. 

According  to  the  method  of  preparation  silicic  acid  may  be 
either  a  negative  or  a  positive  colloid.  When  it  is  a  negative 
colloid  it  retains  some  anion,  such  as  the  chloride  ion,  Cl~,  which 
cannot  be  removed  by  dialysis  or  washing.  When  it  is  a  positive 
colloid  it  retains  some  cation  as  the  sodium  ion,  Na+,  in  the  same 
way.  Solutions  of  the  first  class  are  precipitated  by  solutions 
containing  a  bivalent  cation,  such  as  barium  chloride,  BaCl2, 
while  those  of  the  second  class  are  precipitated  by  solutions 
containing  a  bivalent  anion,  such  as  potassium  sulfate,  K2SO4 
(see  p.  261). 

Colloidal  silicic  acid,  which  is  present  in  arable  soils,  retains 
cations  in  a  form  which  cannot  be  washed  out  by  the  rain, 


354  A  TEXTBOOK  OF  CHEMISTRY 

probably  owing  to  the  relations  which  have  just  been  given,  and 
potassium  seems  to  be  held  much  more  strongly  than  magnesium, 
calcium  or  sodium.  This  is  doubtless  of  great  importance  in  its 
relation  to  the  fertility  of  the  soil,  since  potassium  is  one  of  the 
most  essential  elements  for  the  growth  of  crops. 

If  hydrochloric  acid  is  added  to  the  solution  of  sodium  silicate 
drop  by  drop,  instead  of  suddenly,  the  silicic  acid  separates  as  a 
gelatinous  precipitate.  If  this  is  dried  on  the  water  bath  or  at  a 
slightly  higher  temperature,  or  if  the  colloidal  solution  obtained 
by  the  sudden  acidification  is  evaporated  and  the  residue  dried, 
the  silicic  acid  becomes  almost  completely  insoluble  in  acid 
solutions.  Such  a  process  is  much  used  in  the  quantitative 
analysis  of  silicates. 

The  gelatinous  precipitate  obtained  by  precipitation  has  very 
nearly  the  composition  Si(OH)4  (Norton,  J.  Am.  Chem.  Soc. 
19,  832  (1897)),  but  there  is  considerable  doubt  whether  it  is  a 
definite  compound,  since  it  loses  water  very  easily  and  its  vapor 
pressure  is  scarcely,  if  at  all,  different  from  that  of  pure  water. 
If  dried,  however,  it  loses  the  last  portions  of  water  with  difficulty 
and  only  when  heated  to  bright  redness.  Opal  must  be  con- 
sidered as  a  mixture  of  silicic  acids,  (SiO2.H2O),  but  is  variable 
in  composition,  and  it  cannot  be  said  that  the  existence  of  any 
definite  compound  having  the  composition  of  a  silicic  acid  has 
been  established.  Silicon  dioxide,  SiO2,  resembles  carbon 
dioxide,  CO2,  in  this  respect,  but  with  the  difference  that  while 
carbonic  acid,  H2CO3,  dissociates  directly  into  carbon  dioxide 
and  water  the  hydrates  of  silicon  dioxide  lose  water  gradually, 
forming  hydrates  containing  less  and  less  water,  that  are  united 
together  to  form  complex  molecules,  perhaps  somewhat  as 
follows  : 


=0  ->  H—  Q—  SiO—  O—  SiO—  OH 
OH  H 

The  final  product  of  the  dehydration  of  the  silicic  acid  would, 
according  to  this  view,  consist,  not  of  simple  molecules  of  silicon 
dioxide  but  of  complex  molecules  (SiO2)ra  in  which  the  molecules 


NATURAL  SILICATES  355 

are  held  together  through  the  agency  of  oxygen.  The  large 
number  of  complex  silicates  which  are  known  furnish  a  strong 
support  for  this  view.  The  extremely  high  melting  point  and 
boiling  point  of  silica,  especially  as  compared  with  carbon  di- 
oxide, also  indicate  that  its  molecule  is  complex  and  that  it  does 
not  have  the  simple  formula  SiO2. 

Natural  Silicates.  Carbon  forms  a  great  number  of  acids 
in  which  the  complexity  is  due  to  the  union  of  carbon  atoms  with 
each  other  and  with  varying  numbers  of  other  atoms  and  groups, 
and  it  forms  only  a  single  acid  for  which  carbon  dioxide  is  the 

-0\  /O- 

anhydride.  Such  a  grouping  as  O=C — O— C=O  seems  to 
be  extremely  unstable.  As  has  just  been  pointed  out,  how- 
ever, molecules  containing  silicon  atoms  united  by  oxygen  seem 
to  be  stable,  and  a  great  variety  of  minerals  are  known  which  may 
be  considered  as  salts  of  more  or  less  complex  silicic  acids,  all  of 
which  are  derived  from  the  same  -anhydride,  silica,  SiO2.  The 
hypothetical  acids  from  which  these  natural  silicates  are  derived 
are  best  classified  according  to  the  number  of  atoms  of  silicon 
contained  in  one  molecule  of  the  acid.  The  first  two  are  given 
names  similar  to  the  names  of  the  acids  of  phosphorus,  arsenic 
and  antimony : 

Orthosilicic  Acid  H4SiO4 

Metasilicic  Acid  H2SiO3 

Disilicic        Acid  H6Si2O7 

Trisilicic       Acids          H4Si3O8  and  H8Si3Oi0 
The  following  minerals  may  be  given  as  illustrations  of  the  salts 

of  the  above  acids. 

^    ,     _  Calamine,  Zn2SiO4.H2O 

Orthosmcates ; 

Derivatives  of       ^^  H2Al2(SiO4),H2O 
H4bl°4  Garnet,  Ca3Fe2(SiO4)3 


Metasilicates ; 
Derivatives  of 
H2Si03 


Amphibole 

Hornblende  I CaMg3(SiO3)4 

Asbestos      J 

Talc  or  soaps  tone,  Mg3H2(SiO3) 

Beryl,  Be3Al2(SiO3)6 


356 


A  TEXTBOOK  OF  CHEMISTRY 


Disilicates ; 
Derivatives  of 
H6Si2O7  Serpentine,  I^MgaSiaOc,  (or  Mg3Si2O7.2  H2O) 

Trisilicates ;  . , 

Derivatives  of      ( Orthoclase,  KAlSi3O8 

H4Si308  iAIbite,  NaAlSi3O8 

Derivatives  of 
H8Si3Oi0  Meerschaum,  H4Mg2Si3Oi0 

*  Calculation  of  the  Formula  of  a  Mineral.  The  following 
analysis  of  a  garnet  may  be  taken  as  typical  of  the  analysis  of  a 
silicate : 


PER  CENT 

PER  CENT  OP 
OXYGEN 

RATIOS 

Silica,  SiO2       

40.45 
19.67 
4.05 
2.60 
6.90 
5.78 
20.79 

21.57 

9.26 
1.21 
0.81, 
1.53 
1.65 
8.32 

21.57 
11.28 

11.50 

2  or  6 
lor  3 

1  or  3 

Alumina  A^Oa 

Ferric  oxide,  Fe2Os    

Chromic  oxide,  Cr2Oa    .... 
Ferrous  oxide,  FeO   

Calcium  oxide,  CaO       .... 
Magnesium  Oxide,  MgO    .     .     . 

100.24 

In  this  and  in  similar  analyses  of  minerals  and  rocks  it  is 
customary  to  calculate  the  results  as  though  the  elements  were 
present  as  oxides.  This  custom  was  originally  based  on 
Lavoisier's  system  of  nomenclature,  according  to  which  salts 
were  considered  as  compounds  of  oxides  of  the  metals  with  oxides 
of  the  nonmetals,  and  it  was  continued  during  the  first  half  of 
the  nineteenth  century  largely  because  of  the  older  electrochemi- 
cal theory,  which  seemed  to  give  a  satisfactory  theoretical  basis 
for  the  old  nomenclature.  This  method  of  calculation  is  still 
retained  because  many  of  the  silicates  can  be  prepared  by  the 
direct  union  of  the  oxides  of  the  metals  with  silica,  but  also  for 
the  practical  reason  that  when  a  silicate  analysis  is  calculated 


DIALYSIS,  SEMIPERMEABLE  MEMBRANES      357 


in  this  manner  the  failure  of  the  percentages  found  to  add  up  to 
100  shows  at  once  that  there  is  an  error  in  the  analysis  or  that 
some  element  has  been  overlooked. 

In  calculating  a  formula  for  a  mineral  it  is  convenient  to  make 
use  of  the  principle  that  the  amounts  of  oxygen  in  the  different 
oxides  must  be  in  a  simple  ratio  to  each  other.  Thus  if  we  write 
calcium  silicate  CaOSiO2,  the  ratio  of  the  oxygen  in  the  calcium 
oxide  ("  lime  ")  must  be  to  that  in  the  silicon  dioxide  (silica) 
as  1:2.  An  examination  of  the  quantities  of  oxygen  in  the 
second  column  of  figures  above  shows,  however,  no  simple 
ratio  between  the  amounts  of  oxygen  in  the  various  oxides.  It 
is  not  till  we  put  together  the  oxygen  of  the  oxides  of  ferrous 
iron,  calcium  and  magnesium  and  that  of  the  oxides  of  ferric 
iron,  chromium  and  aluminium  that  we  obtain  numbers  which 
form  simple  ratios.  When  we  do  this  we  obtain,  approximately, 
the  ratios  R"O  :  R2///O3  :  SiO2  =  3:1:3,  and  may  write  the  gen- 
eral formula  of  garnet,  3  R"O.R2'X'O3.3  SiO2  or  R3"R'"(SiO4)3, 
in  which  R"  stands  for  ferrous  iron,  calcium  or  magnesium  and 
R'"  stands  for  ferric  iron,  chromium  or  aluminium. 

If  this  formula  is  written  in  the  ordinary  way,  it  becomes 
3  CaO.Al2O3.3  SiO2  or  Ca3Al2(SiO4)3,  and  we  see  that  garnet  is 
a  derivative  of  orthosilicic  acid,  H4SiO4. 

It  is  fair  to  say,  however,  that  the  published  analyses  of  many 
of  the  complex  silicates  agree  only  very  roughly  with  the  for- 
mulas which  have  been  assigned  to  the  minerals.  The  condi- 
tions under  which  such  minerals  have  been  formed  in  nature  have 
evidently  favored  the  formation  both  of  complex,  isomorphous 
mixtures  and  of  solid  solutions  of  variable  composition. 

Dialysis,  Semipermeable  Membranes.  Colloidal  silicic  acid 
may  be  separated  from  sodium  chloride  by  dialysis,  with  the 
use  of  parchment  or  parchment  paper.  Graham,  who  first 
carefully  studied  phenomena  of  this  kind,  distinguished  between 
crystalloids  and  colloids  in  regard  to  this  effect  of  animal  mem- 
branes. As  he  used  these  names  they  imply  that  crystalline 
compounds,  such  as  sodium  chloride,  most  salts  and  ordinary 
acids  will  pass  through  the  parchment,  while  silicic  acid,  albu- 


358  A  TEXTBOOK  OF  CHEMISTRY 

men  and  other  substances  which  do  not  crystallize,  or  which 
crystallize  with  difficulty,  will  not  pass  through.  While  this  dis- 
tinction still  has  considerable  force,  in  a  general  way,  a  fuller 
knowledge  of  the  subject  of  colloids  has  led  to  a  definition  of 
them  (p.  261)  which  is  based  on  quite  other  properties  than  their 
relation  to  separating  membranes,  and  the  term  crystalloid  is 
now  little  used. 

A  membrane  or  septum  which  allows  one  substance  to  pass 
through  it  while  it  prevents  the  passage  of  another  is  called 
semipermeable.  A  piece  of  parchment  will  allow  water  or  salt  in 
solution  to  pass,  but  is  nearly  impervious  to  silicic  acid  or  albumen. 
It  will  allow  cane  sugar  to  pass,  but  more  slowly  than  salt.  If  a 
precipitate  of  copper  ferrocyanide,  C^FeCeNe,  is  formed  inside 
of  the  wall  of  a  porous  porcelain  cup,  by  placing  a  copper  sulfate, 
CuSO-i,  solution  within  and  a  solution  of  potassium  ferrocyanide, 
TQFeCeNe,  on  the  outside,  the  gelatinous  membrane,  formed  at 
the  point  in  the  wall  where  the  two  solutions  come  together  is  per- 
meable to  water  but  may  be  made  wholly  impervious  to  cane  sugar. 
A  thin  sheet  of  India  rubber  is  readily  permeable  to  pyridine  but 
nearly  impervious  to  glucose  in  solution  in  the  pyridine.  Metal- 
lic palladium  is  permeable  to  hydrogen  but  impermeable  to  ni- 
trogen and  most  other  gases. 

The  mechanism  of  the  permeability  is  probably  different  in 
different  cases.  Palladium  dissolves  hydrogen  very  much  as 
water  dissolves  carbon  dioxide,  but  will  give  it  up  again  to  a 
vacuum  or  to  any  space  containing  no  hydrogen,  although  it 
may  contain  another  gas.  There  is  some  difference  of  opinion 
about  the  permeability  of  copper  ferrocyanide.  Some  authors 
think  the  pores  of  the  membrane  are  so  fine  that  molecules 
of  water  can  pass  through  them  while  the  larger  molecules  of 
sugar  are  stopped.  Others  think  that  the  membrane  dissolves 
water  on  one  side  and  gives  it  out  on  the  other,  as  the  palladium 
dissolves  hydrogen  and  gives  it  up. 

Osmotic  Pressure.  The  passage  of  a  liquid  or  solution 
through  a  membrane  in  the  manner  which  has  been  described  is 
called  osmosis.  If  a  concentrated  solution  of  copper  nitrate  is 


OSMOTIC  PRESSURE 


359 


placed  in  the  parchment  sack,  Fig.  91,  it  will  be  seen  very  soon 
that  the  solution  inside  of  the  sack  is  at  a  higher  level  than  the 
water  on  the  outside,  indicating  a  greater  pressure  on  the  side 
of  the  solution  and  showing  that 
some  water  has  passed  through 
the  membrane  into  the  solution. 
A  pressure  developed  in  this 
manner  is  called  osmotic  pres- 
sure. The  pressures  developed 
with  the  parchment  sack  will  be 
small,  partly  because  the  salt 
solution  as  well  as  the  water 
passes  through  the  membrane, 
which  does  not  form  a  perfect 
septum,  and  partly  because  any 
considerable  difference  of  pres- 
sure between  the  two  sides  would 
burst  the  parchment. 

By  the  arrangement  shown  in 
Fig.  92  it  is  possible  to  measure 
an  osmotic  pressure  of  many 
atmospheres.  A  membrane  of 
copper  ferrocyanide,  C^FeCeNe, 
is  first  prepared,  as  described  in 
the  last  paragraph,  within  the 
walls  of  the  porous  porcelain 
cup,  z.  The  tube,  m,  is  filled 
with  mercury,  leaving  air  in  the 
graduated  capillary  tube.  The 
rest  of  the  apparatus  is  then 
completely  filled  with  a  solution 
of  sugar  or  of  some  other  sub- 


Fig.  92 


stance  which  is  to  be  examined,  and  the  cup,  z,  is  placed  in 
distilled  water.  The  volume  of  air  in  the  capillary  tube  will 
soon  begin  to  diminish,  and  the  contraction  will  continue  till  a 
pressure  on  the  solution  is  developed  which  will  just  prevent  the 


360  A  TEXTBOOK  OF  CHEMISTRY 

further  passage  of  water  through  the  membrane ;  the  system  is 
then  in  equilibrium.  This  hydrostatic  pressure  on  the  solution, 
necessary  to  prevent  the  passage  of  the  solvent  through  a  semi- 
permeable  membrane  into  a  solution  is  the  osmotic  pressure  of 
the  solution  (van't  Hoff).  It  is  evident  that  if  equilibrium  is 
reached  when  the  volume  of  the  air  in  the  tube  is  one  half  the 
original  volume,  the  pressure  of  the  air  must  be  two  atmospheres 
and  the  osmotic  pressure  must  be  the  difference  between  this 
pressure  and  the  pressure  of  the  air  on  the  water  outside,  which 
would  be  one  atmosphere.  If  the  air  in  the  tube  is  reduced  to 
one  third  its  original  volume,  the  osmotic  pressure  must  be  two 
atmospheres.  * 

The  osmotic  pressures  developed  in  this  manner  are  very  con- 
siderable. A  tenth-formular  solution  of  cane  sugar,  C^H^On 
(containing  34.2  grams  in  a  liter  or  3.42  per  cent),  will  give  an 
osmotic  pressure  at  0°  of  about  2.24  atmospheres,  while  a  formu- 
lar  solution  of  alcohol,  C2H5OH  (containing  4.6  per  cent  by 
weight),  would  give  a  pressure  of  more  than  20  atmospheres.  The 
pressures  are  very  nearly  proportional  to  the  absolute  tempera- 
ture and,  in  dilute  solutions,  nearly  the  same  as  though  the  dis- 
solved substance  (solute)  were  in  the  form  of  a  gas  in  the  same 
volume  and  at  the  same  temperature. 

It  will  be  seen,  at  once,  that  a  measurement  of  the  osmotic 
pressure  may  be  used  to  determine  the  molecular  weight  of  a 
compound  exactly  as  the  density  of  a  gas  or  vapor  is  used  for  this 
purpose  (p.  94).  It  has  been  shown,  also,  by  processes  of  reason- 
ing which  it  would  carry  us  too  far  to  give  here,  that  there  is  a 
necessary  connection  between  the  osmotic  pressure  of  a  solution 
and  its  vapor  pressure,  boiling  point  and  freezing  point  (van't 
Hoff).  The  use  of  the  freezing  point  of  solutions  to  determine 
molecular  weights  (p.  112)  is  intimately  connected  with  these 
relations. 

A  mental  picture  of  what  may  be  the  cause  of  osmotic  pressure 
can  be  formed  by  considering  the  effect  of  a  septum  of  palladium 
on  hydrogen  gas.  If  a  bulb  of  palladium  containing  nitrogen 
under  atmospheric  pressure  is  placed  in  an  atmosphere  of  hydro- 


GERMANIUM  361 

gen,  also  under  atmospheric  pressure,  hydrogen  will  be  absorbed 
by  the  palladium  on  the  outside  and  given  off  on  the  inside  until 
the  pressure  of  the  hydrogen  on  the  two  sides  of  the  wall  of 
palladium  is  the  same.  The  pressure  on  the  inside  of  the  bulb 
will  be  two  atmospheres,  one  atmosphere  due  to  nitrogen  and 
one  atmosphere  due  to  hydrogen.  This  is  because  nitrogen  gas 
is  discontinuous  and  offers  no  resistance  to  the  escape  of  mole- 
cules of  hydrogen  from  the  surface  of  the  palladium.  When 
inside,  the  nitrogen  and  hydrogen  each  exert  their  normal  pres- 
sure on  the  wall  as  gases. 

In  a  similar  manner,  if  we  have  a  dilute  solution  of  sugar  on 
one  side  of  a  membrane  of  copper  ferrocyanide  and  pure  water 
on  the  other  side,  the  water  will  pass  through  till  the  pressure 
due  to  the  water  is  the  same  on  both  sides.  When  equilibrium 
is  reached  the  pressure  on  the  side  of  the  sugar  solution  will  be 
greater  than  that  on  the  side  of  the  pure  water  by  the  amount 
of  the  pressure  due  to  the  sugar.  It  has  been  shown  experi- 
mentally that  this  pressure  is  very  nearly  the  same  as  though 
the  solute  existed  as  a  gas  in  the  same  space. 

It  is  clear  from  what  has  been  said  that  the  kinetic  theory 
(p.  58)  must  apply  in  a  modified  form  to  liquids  as  well  as  gases. 
We  shall  find  later  (p.  397)  that  it  applies  to  solids  also,  as  shown 
by  the  specific  heats  of  the  elements. 

*  Germanium,  Ge,  72.5.  In  1886  Clemens  Winkler  found, 
after  repeated  analyses  of  a  mineral  called  argyrodite,  that  the 
sum  of  the  elements  found  was  always  6  to  7  per  cent  less  than 
100.  This  led  him  to  a  further  careful  study  of  the  material 
and  to  the  discovery  of  germanium,  an  element  intermediate 
in  its  properties  between  silicon  and  tin,  but  on  the  whole  resem- 
bling the  latter  much  more  closely  than  the  former.  It  is  a 
brittle,  grayish  white  metal,  which  melts  at  958°.  It  forms  the 
compounds  GeO,  GeS  and  GeCl2,  in  which  it  is  bivalent  and 
GeH4,  GeHCl3,  GeCl4,  GeO2,  GeS2  and  K2GeF6,  all  of  which 
correspond  to  similar  compounds  of  quadrivalent  silicon. 

Argyrodite  has  approximately  the  composition  4  Ag2S.GeS2. 

Tin,  Sn,  119,  and  Lead,  Pb,  207.1,  are  very  distinctly  metallic 


362  A  TEXTBOOK   OF  CHEMISTRY 

in  their  properties  in  the  free  state  and  also  in  the  formation  of 
salts  in  which  they  are  the  metallic  elements.  They  give,  how- 
ever, the  oxides  SnO2,  stannic  oxide,  which  is  the  anhydride  of 
stannic  (H2SnO3),  and  metastannic  acids,  and  PbO2,  lead  dioxide, 
which  is  the  anhydride  of  plumbic  acid,  H4PbO4.  These  metals 
and  their  compounds  will  be  considered  more  in  detail  later. 

The  elements  of  the  fourth  group  of  the  periodic  system, 
which  are  found  in  the  alternate  rows,  show  a  closer  resemblance 
to  silicon  and  germanium  than  is  shown  by  tin  and  lead.  It 
seems,  on  the  whole,  therefore,  better  to  give  an  account  of  them 
here  rather  than  among  the  metallic  elements. 

*  Titanium,  Ti,  48. 1 .  This  element  is  found  in  the  mineral  rutile 
TiO2,  in  many  iron  ores,  especially  in  the  magnetic  iron  ore,  and 
in  small  amounts  in  almost  all  rocks.  Ordinary  soils  contain,  on 
the  average,  more  than  half  a  per  cent  of  titanium  oxide  and  it  is 
probably  a  constituent  of  the  ash  of  most  plants  and  animals. 
The  free  element  has  apparently  never  been  obtained  pure, 
largely  because  of  its  strong  affinity  for  nitrogen  and  carbon. 
The  purest  specimens  which  have  been  obtained  have  a  very  high 
melting  point,  1790°,  suggesting  the  possibility  of  using  the  ele- 
ment for  the  filaments  of  electric  lamps,  but  no  one  has  succeeded 
in  applying  it  to  this  purpose.  Titanium  forms  compounds 
in  which  the  element  is  bivalent  (TiCl2,  TiO,  TiS),  trivalent 
(TiCl3,  Ti2O3,  Ti2S3,  Ti2(SO4)3,  TiN)  and  quadrivalent  (TiCl4, 
Ti02,  TiS2,  K2TiF6). 

Titanium  tetrafluoride,  TiF4,  is  a  white  powder  which  boils 
at  284°,  but  it  is  much  more  easily  decomposed  by  water  than 
silicon  tetrafluoride  so  that  on  treatment  of  a  mixture  of  silicon 
dioxide,  SiO2,  and  titanium  dioxide,  TiO2,  with  a  mixture  of 
hydrofluoric  and  sulfuric  acids,  the  silicon  may  be  expelled  as 
the  tetrafluoride  on  evaporation,  while  the  titanium  remains  be- 
hind as  the  oxide.  Acid  solutions  containing  titanium  give  a 
deep  yellow  color  with  hydrogen  peroxide,  due  to  the  formation 
of  pertitanic  acid.  This  is  used  for  the  detection  and  estimation 
of  titanium  and  may  also  be  used  for  the  detection  of  hydrogen 
peroxide.  Hydrofluoric  acid  interferes  with  the  reaction. 


CERIUM  363 

Some  compounds  of  titanium  are  used  as  mordants  in  dyeing 
wool,  cotton  and  leather. 

*  Zirconium,  Zr,  90.6,  is  found  most  frequently  in  the  mineral 
zircon,  ZrSiO4.  It  is  also  found  in  considerable  quantities  in 
Brazil  in  the  form  of  the  dioxide,  ZrO2.  The  element  exists 
in  two  forms.  The  amorphous  form  is  a  black  powder,  resem- 
bling carbon.  The  crystalline  form  is  brittle,  melts  at  about  1700° 
and  has  a  specific  gravity  of  6.4.  Zirconium  hydride,  ZrH2,  is 
a  black  powder.  The  chloride,  ZrCU,  is  hydrolyzed  by  water. 
Zirconates,  which  resemble  some  of  the  silicates,  are  prepared  by 
fusing  zircon  dioxide,  ZrO2,  with  metallic  oxides,  carbonates  or 
chlorides.  Crystalline  silicozirconates  have  also  been  prepared. 

Zirconium  dioxide  glows  very  intensely  in  the  oxyhydrogen 
flame,  giving  an  even  better  light  than  lime,  and  the  light  is  used 
for  intensive  illumination  in  spectroscopy  and  microphotography. 
The  oxide  is  used  with  yttrium  oxide  as  the  chief  constituent  of 
the  Nernst  lamp.  The  mass  conducts  electricity  well  only  at 
high  temperatures  and  must  be  heated  in  some  way  to  start 
the  passage  of  the  current.  It  was  at  one  time  used  as  the  incan- 
descent material  for  the  Welsbach  light,  but  that  use  has  been 
abandoned.  There  seems  to  be  some  possibility  of  using  zircon 

tor  some  of  its  compounds  in  the  filaments  of  electric  lamps. 
Fused    zirconium  dioxide  resembles   fused  quartz   and  cru- 
cibles for  laboratory  use  have  been  made  from  it. 

Cerium,  Ce,  140.25.  In  Brazil  and  in  North  and  South 
Carolina  there  are  found  large  deposits  of  a  heavy  sand,  called 
monazite  sand,  consisting  of  a  complex  mixture  of  heavy, 
insoluble  minerals  which  have  been  sorted  out  from  other 
materials,  partly  by  the  disintegration  of  rocks  in  which  they 
were  originally  disseminated,  partly  by  a  process  of  washing 
away  the  lighter  minerals  in  operations  which  doubtless  occurred 
during  very  long  periods  of  geological  time.  These  sands  con- 
tain a  great  variety  of  minerals  and  are  of  considerable  commer- 
cial value  because  of  the  thorium  and  cerium  in  them.  The  ce- 
rium seems  to  be  mostly  present  as  cerium  phosphate,  CePO4. 

Metallic  cerium  is  an  iron-gray,  malleable  metal,  which  burns 


364  A  TEXTBOOK  OF  CHEMISTRY 

easily  to  cerium  dioxide,  CeO2,  in  the  air.  Its  specific  gravity 
is  6.73.  An  alloy  with  iron  throws  off  incandescent,  burning 
particles  on  striking  with  a  piece  of  steel.  These  will  ignite  gas 
or  alcohol  and  the  alloy  is  used  for  gas  lighters  and  for  similar 
purposes. 

*  Cerium  forms  three  oxides,  Ce2O3,  CeO2  and  CeOs.  The 
dioxide,  CeO2,  forms  1  per  cent  of  the  material  used  for  the 
Welsbach  mantles.  (See  below.)  Cerous  sulfate,  Ce2(SC>4)3, 
forms  a  difficultly  soluble  double  salt  with  sodium  sulfate, 
Ce2(SO4)s.Na2SO4.2  H2O,  which  is  used  in  separating  cerium 
from  other  minerals.  The  eerie  sulfate,  Ce(SO4)2.4  H2O,  in 
which  the  cerium  is  quadrivalent  is  readily  hydrolyzed  to  basic 
sulf  ates . 

Thorium,  Th,  232.4,  is  found  in  monazite  sand,  probably  as 
the  dioxide,  ThO2,  and  this  forms  the  practical  source  for  thorium 
compounds,  which  have  become  very  important  for  use  in  Wels- 
bach mantles.  Thorium  is  also  found  in  the  mineral  thorianite, 
an  isomorphous  mixture  of  thorium  dioxide,  ThO2,  and  uranium 
oxide,  UO2.  Thorianite  is  found  in  Ceylon.  The  mineral  has 
proved  of  unusual  interest  as  the  source  of  radiothorium,  one  of 
the  strongly  radioactive  elements. 

Apparently  pure  metallic  thorium  has  not  been  prepared. 
From  the  properties  of  the  impure  metal  it  probably  has  a  spe- 
cific gravity  of  about  12.2  and  melts  at  1700°  or  above. 

Thorium  forms  many  salts  in  which  it  is  quadrivalent,  of  which 
the  sulfate,  Th(SO4)2.9  H2O,  and  the  nitrate,  Th(NO3)4.12  H2O, 
may  be  considered  as  typical. 

Welsbach  Mantles.  Compounds  of  thorium  and  cerium  are 
now  important  articles  of  commerce  for  the  manufacture  of  the 
mantles  used  for  the  Welsbach  light.  In  the  preparation  of 
these  mantles  a  web  of  cotton  or,  more  recently  and  better,  of 
artificial  silk  is  dipped  in  a  solution  containing  nitrates  of  cerium 
and  thorium.  Careful  experiments  have  shown  that  the  best 
results  are  obtained  with  a  mantle  containing  99  per  cent  of 
thorium  dioxide,  ThO2,  and  1  per  cent  of  cerium  dioxide,  CeO2. 
After  drying,  the  cotton  or  artificial  silk  is  burned  out  and  the 


BORON  365 

skeleton  of  oxides  is  dipped  in  collodion  and  dried.  The  elastic 
film  of  nitrocellulose  left  by  the  evaporation  of  the  collodion 
holds  the  fragile  skeleton  of  oxides  together  for  transportation. 

A  study  of  the  spectra  of  the  light  from  mantles  made  of  pure 
thorium  oxide,  pure  cerium  oxide  and  of  mixtures  of  the  two  has 
shown  that  thorium  oxide  radiates  comparatively  little  light  of 
the  wave  lengths  of  the  visible  spectrum,  apparently  because  it 
is  nearly  transparent.  Cerium  oxide,  on  the  other  hand, 
radiates  so  much  energy  of  the  wave  length  of  the  ultra  red  rays 
that  it  lowers  the  temperature  of  the  flame  too  far  to  produce  a 
satisfactory  light.  In  the  mixture,  the  thorium  dioxide,  which 
has  a  very  low  specific  heat  and  slight  power  of  emission,  is 
heated  to  a  high  temperature  (1500°-1600°),  and  at  this  tempera- 
ture the  cerium  oxide  has  a  greatly  increased  power  of  radiation 
for  rays  of  the  wave  lengths  of  the  visible  spectrum.  Small 
amounts  of  chromium,  platinum,  manganese  or  uranium  may 
produce  effects  similar  to  those  produced  by  the  cerium,  but  none 
of  these  give  as  successful  a  mixture  as  that  of  the  oxides  of 
thorium  and  cerium. 

Boron,  B,  11.  In  some  places  in  Tuscany,  in  northern  Italy, 
steam  which  is  charged  with  boric  acid,  HaBOs,  escapes  from 
subterranean  sources.  The  boric  acid  is  obtained  by  bringing 
the  steam  into  contact  with  water,  which  absorbs  the  acid  and  by 
evaporating  the  solution  till  it  crystallizes  on  cooling.  Borax, 
Na2B4O7.10  H2O,  is  found  in  many  natural  waters  and  especially 
in  Borax  Lake,  in  California.  A  calcium  salt,  called  colemanite, 
Ca2B6On.5  H2O,  is  also  found  in  veins  in  California  and  is  used 
for  the  manufacture  of  borax. 

Preparation,  Properties.  Amorphous  boron  may  be  prepared 
by  heating  the  trioxide,  B2O3,  or  dehydrated  borax,  Na2B4O7, 
with  magnesium.  It  forms  a  brownish  gray  powder.  Boron 
may  be  obtained  in  a  crystalline  form  by  crystallizing  it  from 
metallic  aluminium.  The  crystals  are  extremely  hard,  approach- 
ing the  hardness  of  the  diamond.  Boron  melts  at  2200°-2500°. 

Boron  Trioxide,  B2O3,  Borax  Beads.  The  trioxide  or  boric 
anhydride,  B2Oa,  is  formed  by  heating  either  of  the  oxygen  acids 


366  A  TEXTBOOK  OF  CHEMISTRY 

of  boron  to  a  high  temperature.  It  is  volatile  only  at  a  white 
heat  and  on  this  account  it  will  decompose  the  salts  of  almost 
all  other  acids,  when  heated  with  them  to  a  high  temperature. 
It  does  this  in  spite  of  the  fact  that  in  solution  boric  acid  is  one 
of  the  weakest  acids  known — so  weak,  in  fact,  that  its  solution 
does  not  taste  sour  and  does  not  redden  litmus.  When  fused 
with  salts  or  with  the  oxides  of  metals,  boric  anhydride  combines 
with  the  oxides  to  form  borates,  just  as  silicon  dioxide  combines 
with  oxides  to  form  silicates.  Borax  glass,  Na2B4O7,  which  may 
be  considered  as  boric  anhydride  that  has  been  only  partly 
neutralized  by  sodium  oxide  (Na2O.2  B2O3),  decomposes  salts 
or  combines  with  oxides  of  metals  in  the  same  manner,  and 
almost  all  of  the  borates  of  metals  prepared  in  this  way  remain 
dissolved  in  the  anhydride  or  combined  with  it,  giving  a  mixture 
of  borates,  which  solidifies  to  a  clear  glass  on  cooling.  Many 
of  these  glasses  have  characteristic  colors  and  so  the  borax  bead 
is  much  used  for  the  detection  of  metals  in  mineralogy  and 
qualitative  analysis. 

Some  of  the  colorless  borates  may  be  fused  with  silicates  to 
form  clear,  transparent  glasses,  which  are  much  less  soluble  than 
ordinary  glass.  Such  borosilicates  are  now  used  in  the  manu- 
facture of  glassware  for  chemical  laboratories. 

Boric  Acid,  HsBOs.  If  sulfuric  acid  or  hydrochloric  acid  is 
added  to  a  moderately  concentrated  solution  of  borax, 
Na2B4O7.10  H2O,  in  an  amount  equivalent  to  the  sodium 
present,  orthoboric  acid,  H3BO3,  crystallizes  from  the  solution 
on  cooling.  It  is  usually  called  simply  boric  acid  and  is  a  very 
weak  acid,  having  no  sour  taste  and  none  of  the  corrosive 
qualities  of  ordinary  acids.  It  dissolves  in  about  25  parts  of 
water  at  15°.  The  solution  has  powerful  germicidal  properties 
and  is  often  used  for  an  eyewash.  Both  boric  acid  and  borax 
have  been  used  as  food  preservatives,  but  such  a  use  is  con- 
demned because  the  substances  produce  poisonous  effects  when 
taken  in  considerable  quantities. 

Other  Acids  of  Boron.  At  100°  boric  acid  loses  one  molecule 
of  water  and  is  changed  to  metaboric  acid,  HBO2,  and  at  140° 


OTHER  COMPOUNDS  OF  BORON  367 


it  loses  still  more  water  and  is  converted  into  pyroboric  acid, 
H2B407. 

Borax,  Na2B4O7.10  H2O,  is  usually  considered  as  a  salt  of 
pyroboric  acid,  H2B4O7,  but  it  may  with  equal  propriety  be 
called  an  acid  salt  of  boric  acid  and  the  formula  written 
Na2Hio(BO3)4.5  H2O.  When  heated  it  loses  water  and  finally 
melts  to  a  clear  glass  consisting  of  sodium  pyroborate,  Na2B4O7. 
The  conduct  of  this  glass  toward  metallic  oxides  and  its  use  in 
blowpipe  analysis  have  been  mentioned  above.  Because  of 
these  properties  it  is  used  as  a  flux  in  assaying  and  in  metallurgy. 
It  is  sometimes  sprinkled  over  iron  which  is  to  be  welded.  In 
this  use  it  dissolves  the  oxide  of  iron  on  the  surface,  forming  an 
easily  fusible  glass  which  is  pushed  out  from  between  the  two 
surfaces  when  the  joint  is  hammered,  leaving  clean  surfaces  of 
iron  to  unite  in  the  weld. 

Borax  is  used  for  laundry  purposes,  having  the  properties  of  a 
mild  alkali.  Its  use  as  an  antiseptic  in  milk  and  other  foods  is 
forbidden  in  most  countries. 

Sodium  Perborate,  NaBO3.4  H2O,  has  recently  come  into  use 
in  laundries  because  it  combines  the  properties  of  a  mild  alkali 
with  the  bleaching  properties  of  hydrogen  peroxide.  Its  struc- 
ture, from  the  methods  of  preparation,  must  be  closely  related 
to  that  of  hydrogen  peroxide,  H — O — O — H,  and  is  probably 
Na— O— O— B=0. 

Other  Compounds  of  Boron.  Boron  forms  a  nitride,  BN,  a 
chloride,  BC13,  and  a  sulfide,  B2S3,  all  of  which  are  hydrolyzed 
by  water.  The  fluoride,  BF3,  gives  boric  acid  and  fluoboric 
acid  by  hydrolysis,  resembling  silicon  fluoride,  SiF4,  in  this 
respect : 

BF3  +  3  HOH  =  H3BO3  +  3  HF 
HF  +  BF3  =  HBF4 

Fluoboric 
Acid 

If  boric  acid  or  a  borate  is  warmed  with  alcohol  and  con- 
centrated sulfuric  acid,  an  ester  of  boric  acid,  (C2H5)3BO3,  is 


368  A  TEXTBOOK  OF  CHEMISTRY 

formed.    This  is  volatile  and  gives  a  green  color  to  the  alcohol 
flame,  as  it  burns. 

If  a  piece  of  turmeric  paper  is  dipped  in  a  solution  of  boric 
acid  or  of  a  borate  which  has  been  made  faintly  acid  with  hydro- 
chloric acid,  the  paper  assumes  a  very  characteristic  red  color 
on  drying. 

EXERCISES 

1.  How  many  grams  of  magnesium  should  be  mixed  with  3  grams  of 
sand  to  form  silicon  ?    How  many  grams  to  form  magnesium  silicide  ? 

2.  In  the  first  case  above  how  many  cubic  centimeters  of  hydro- 
chloric acid,  sp.  gr.  1.10,  containing  20  per  cent  of  the  acid  will  be 
required  to  dissolve  the  magnesium  oxide  formed  ? 

3.  How  many  grams  of  fluorspar,  of  concentrated  sulfuric  acid  and 
of  sand  will  be  required  to  prepare  one  liter  of  a  10  per  cent  solution  of 
fluosilicic  acid,  assuming  that  one  third  of  the  materials  used  fails  to 
react?     The  specific  gravity  of  such  a  solution  is  1.083. 

4.  What  is  the  formula  of  a  mineral  having  the  following  com- 
position ? 

PER  CENT 

Silica,  SiO2 •>    .    .      66.21 

Alumina,  A12O3 19.16 

Potassium  oxide,  K2O 7.38 

Sodium  oxide,  Na2O 7.25 

100.00 

5.  How  much  sodium  carbonate  will  be  required  to  convert  60  grams 
of  quartz  into  sodium  metasilicate  ?     How  many  grams  of  calcium 
carbonate  to  convert  it  into  calcium  metasilicate  ?     How  many  grams 
of  each  to  form  the  orthosilicates  ?     How  many  grams  to  form  a  salt 
of  trisilicic  acid,  HiSisOs  ? 

6.  What  are  the  equations  for  the  reactions  for  the  preparation  of  the 
ethyl  ester  of  boric  acid,  (C2H5)3BO3,  from  borax  ?    What  compounds 
will  be  formed  on  burning  the  ester  ? 

7.  When  boron  nitride  is  heated  to  200°  in  a  current  of  steam,  pyro- 
boric  acid  and  ammonia  are  formed.     What  is  the  equation  for  the 
reaction  ? 

8.  How  much  crystallized  boric  acid  can  be  obtained  from  100  grams 
of  borax  ? 


CHAPTER  XXII 

METALLIC   ELEMENTS.     DIFFERENCES    BETWEEN   METALS 
AND   NON-METALS.     PREPARATION    OF    COMPOUNDS 

Metals  and  Non-metals.  The  general,  physical  differences 
between  the  metals  and  non-metals  are  familiar  to  every  one. 
Iron,  copper  and  gold  may  be  taken  as  typical  of  the  first  class, 
sulfur,  phosphorus  and  carbon  of  the  second.  The  metals  are 
opaque,  except  in  excessively  thin  layers,  are  malleable  and  ductile 
and  good  conductors  of  heat  and  of  electricity.  The  non-metals 
are  transparent  or  translucent,  brittle,  poor  conductors  of  heat 
and  nonconductors  or  poor  conductors  of  electricity.  It  is 
true  that  these  differences,  which  are  so  marked  in  the  elements 
spoken  of  above  as  typical  of  the  two  classes,  are  not  equally 
marked  in  all  cases  and  that  there  is  a  gradation  in  these  proper- 
ties such  that  some  elements  stand  in  a  borderland  between  the 
two,  but  the  physical  properties  mentioned  are  clearly  those  of 
metals  on  the  one  hand  and  of  non-metals  on  the  other. 

In  chemical  properties  the  metals  and  non-metals  show  equally 
strong  contrasts.  In  these  properties  the  most  typical  metals 
may  be  considered  as  sodium  and  potassium  and  the  most 
typical  non-metals  as  fluorine  and  chlorine.  The  chlorides  and 
even  the  hydroxides  of  the  former  ionize  in  solution  in  such  a 
manner  that  the  metal  becomes  the  positive  ion.  The  chlorides 
of  the  nonmetallic  elements,  on  the  other  hand,  are  hydrolyzed 
by  water  with  the  formation  of  hydrochloric  acid  and  an  acid 
containing  the  non-metal.  The  compounds  of  the  non-metals 
with  oxygen  and  hydrogen  ionize  in  solution  with  the  formation 
of  hydrogen  ions  and  some  complex  group  containing  the  non- 
metallic  element.  As  is  the  case  with  the  physical  properties, 
however,  the  chemical  properties  of  the  elements  show  continu- 
ous gradations  from  those  of  the  metals  to  those  of  the  non- 
369 


370  A  TEXTBOOK  OF    CHEMISTRY 

metals.  Some  of  the  chlorides  of  the  metals  are  hydrolyzed  by 
water  and  some  halides  of  elements  which  are  nonmetallic  in 
most  of  their  properties  may  be  formed  in  the  presence  of  water. 
Further  than  this,  some  elements  as  they  combine  with  more  and 
more  oxygen  may  pass  from  metallic  to  distinctly  nonmetallic 
properties,  while  nitrogen  or  even  sulfur,  when  combined  with 
hydrogen  or  hydrocarbon  radicals,  may  form  radicals,  such  as 
NH4  or  (0113)38,  whose  hydroxides  are  bases. 

The  most  typical  of  the  metallic  elements  are  univalent  toward 
chlorine  and  toward  oxygen  and  form  no  compounds  of  any  kind 
in  which  it  is  certain  that  they  have  a  higher  valence.  The  most 
typical  nonmetallic  elements  on  the  other  hand  show  a  varying 
valence,  especially  toward  oxygen. 

The  development  of  the  electron  theory  makes  it  seem  possible, 
or  perhaps  we  may  even  say,  probable,  that  both  the  physical 
and  chemical  properties  which  constitute  the  differences  between 
metals  and  non-metals  are  very  largely  occasioned  by  the  dif- 
ferences in  the  conduct  of  the  atoms  of  the  elements  toward 
electrons.  The  passage  of  electricity  through  a  metallic  con- 
ductor consists  in  the  flow  of  a  stream  of  electrons,  and  metals 
are  supposed  to  be  good  conductors  because  the  electrons  pass 
easily  from  one  atom  to  another  throughout  the  mass.  For 
almost  the  same  reason  the  atom  of  the  metal  becomes  a  posi- 
tive ion  in  solution  because  it  easily  loses  one  or  more  electrons 
—  the  electrons  being  units  of  negative  electricity.  Metals  are 
also  good  conductors  of  heat  because  the  rapidly  moving  electrons 
transfer  energy  from  one  atom  to  another  throughout  the  mass, 
and  they  are  opaque  because  the  electrons  and  atoms  absorb 
and  diffuse  the  light  vibrations  falling  on  the  surface. 

Classification  of  the  Metals.  As  in  the  case  of  the  non-metals, 
the  periodic  system  furnishes  the  most  satisfactory  classification 
for  the  metals.  The  relationships  brought  out  in  this  way  are  not 
always  so  close  as  might  be  desired,  and  there  are  even  sharper 
contrasts  between  the  alternate  metals  of  the  first  group  than 
those  which  have  been  noticed  among  the  elements  of  the  sixth 
and  seventh  groups. 


METALLIC   ELEMENTS 


371 


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372  A  TEXTBOOK  OF  CHEMISTRY 

All  of  the  elements  sufficiently  well  known  so  that  the  atomic 
weights  are  given  in  the  International  Table  of  Atomic  Weights 
are  included  in  the  table,  with  the  exception  of  hydrogen.  The 
nonmetallic  elements  are  inclosed  in  brackets. 

An  inspection  of  the  table  shows  that  even  if  the  noble  gases 
are  counted  as  nonmetallic  elements  more  than  three  fourths 
of  the  elements  known  are  metallic  in  character.  In  spite  of  this, 
the  chemistry  of  the  metallic  elements  offers  less  variety  than  that 
of  the  non-metals  —  largely  because  the  atoms  of  the  metals 
show  less  tendency  to  combine  with  each  other,  or  indeed  with 
other  elements  to  form  complex  radicals.  Metals  usually  sep- 
arate by  themselves  as  positive  ions  in  solution,  while  the  non- 
metals  more  often  form  complex  ions,  such  as  NOs,  SO4  or  ClOs. 

Melting  Points  of  the  Elements.  The  following  table  gives 
all  melting  points  of  the  elements  which  are  known  with  some 
degree  of  accuracy.  The  melting  points  of  the  elements  whose 
names  are  printed  in  capitals  have  been  very  carefully  deter- 
mined, and  are  used  as  standard  temperatures  for  calibrating 
thermometers  and  pyrometers.  As  nearly  as  may  be,  all  values, 
in  particular  the  standard  points,  have  been  reduced  to  a  common 
scale,  the  thermodynamic  scale. 

Preparation  of  Chemical  Compounds.  It  is  clear  from  many 
illustrations  which  have  been  given  in  the  preceeding  pages  that 
many  chemical  reactions  result  in  an  equilibrium  between 
reacting  compounds  such  that  action  seems  to  cease  only  because 
two  opposing  reactions  proceed  with  equal  velocity  in  opposite 
directions.  From  a  theoretical  standpoint  we  are  probably 
justified  in  considering  all  reactions  as  reversible,  but  in  some  of 
them  the  velocities  of  the  reactions  in  opposite  directions  differ 
so  greatly  that  the  equilibrium  lies  very  far  on  one  side  and  the 
reaction  is  practically  complete.  Thus  in  the  very  simple  case 
of  the  reaction  between  hydrogen  and  oxygen : 

2  H2  +  O2  ^±  2  H2O 

at  1000°  when  equilibrium  is  reached,  only  3  parts  in  10,000,000 
of  the  oxygen  and  hydrogen  will  remain  in  the  form  of  the  free 


MELTING  POINTS  OF  THE   ELEMENTS 


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374  A  TEXTBOOK  OF  CHEMISTRY 

gases  (p.  61).  At  ordinary  temperatures  the  velocity  of  com- 
bination and  dissociation  becomes  so  slow,  in  this  case,  that 
either  a  mixture  of  oxygen  and  hydrogen  or  water  may  remain 
apparently  unchanged  for  an  indefinite  length  of  time,  but  we 
know  that  such  a  condition  of  apparent  equilibrium,  dependent 
on  the  slow  velocity  of  a  reaction,  is  a  state  of  instability.  Such 
conditions  are,  however,  very  common,  and  it  has  been  pointed 
out  that  the  varying  velocity  of  different  possible  reactions  is 
of  very  great  practical  importance  in  organic  compounds. 

A  very  large  number  of  the  compounds  of  the  metals  are 
electrolytes,  and  many  of  these  can  be  prepared  in  aqueous  solu- 
tions. The  velocity  of  ionization  reactions  is  so  great  for  all 
ordinary  electrolytes  and  the  interaction  between  ions  in  solu- 
tion is  so  rapid  that  equilibrium  is  reached  almost  instantane- 
ously. The  amount  of  each  substance  present  in  this  equilib- 
rium is  often  influenced  by  some  property  of  one  of  the  substances 
which  removes  it  from  interaction  with  the  others.  Such  prop- 
erties are,  especially,  volatility,  solubility,  degree  of  ionization 
and  the  formation  of  complex  ions. 

Effect  of  Volatility.  If  concentrated  sulfuric  acid  is  dropped 
into  a  concentrated  solution  of  salt,  there  is,  at  first,  no  apparent 
action.  In  the  solution  the  following  reactions  will  be  almost 
instantly  in  a  state  of  equilibrium  : 

NaCl    ±  Na+  +  CT 


H+  +  cr  ^±  HCI 

Na+  +  HS04-  ^±  NaHS04 

It  is  to  be  noticed  that  the  reaction  as  usually  given 
NaCl  +  H2S04  =  HCI  +  NaHSO4 

is  merely  the  final  result  of  intermediate  ionization  reactions  and 
probably  occurs,  at  first,  in  very  insignificant  amounts. 

As  more  sulfuric  acid  is  added,  a  point  will  be  reached  when 
the  solution  becomes  saturated  with  hydrochloric  acid.  The 
addition  of  a  further  amount  of  sulfuric  acid  must  now  result 


VOLATILE   COMPOUNDS  375 

in  the  escape  of  hydrochloric  acid  in  the  gaseous  form,  as  mole- 
cules of  the  compound,  HC1.  This  will  disturb  the  equilibrium 
of  the  reaction : 

H+  +  cr  ^±  HCI 

in  such  a  manner  that  more  of  the  hydrogen  and  chloride  ions 
will  unite  to  form  un-ionized  hydrochloric  acid,  HCI.  This,  in 
turn,  will  cause  the  formation  of  an  increased  number  of  sodium 
and  chloride  ions  from  the  salt,  NaCl,  and  of  hydrogen  and  hy- 
drosulfate  ions,  HSC>4~,  from  the  sulfuric  acid.  Of  course,  the 
more  concentrated  the  original  solution  of  salt  is,  the  more  hy- 
drochloric acid  can  be  obtained  in  gaseous  form,  and  if  we  start 
with  a  saturated  solution  and  a  large  excess  of  solid  salt,  the  latter 
will  pass  into  solution  as  the  reaction  proceeds  and  nearly  all  of 
the  chlorine  may  escape  as  hydrochloric  acid.  This  result 
occurs  because  the  hydrochloric  acid  is  a  gas  and  escapes  from  the 
mixture  as  a  gaseous  phase  and  in  spite  of  the  fact  that  the 
ionization  reactions  of  hydrochloric  and  of  sulfuric  acids  result 
in  the  formation  of  more  of  the  sulfuric  than  of  the  hydrochloric 
acid  when  the  two  acids  are  present  in  equivalent  amounts. 

When  a  solution  of  hydrochloric  acid  is  dropped  into  a  solu- 
tion of  sodium  carbonate  the  following  reactions  occur : 

HCI  ^±  H+  +  Cl- 
Na2C03  ^±  Na+  +  Na+  +  CO3= 
CO3=  +  H^  ^±  HCOr 
H+  +  HCOr  ^±  H2C03 

H2C03  ;£  H20  +  CO2 
Na+  +  Cl~  ^±  NaCl 
Na+  +  HCOr  ^±  NaHC03 

In  the  third  of  these  reactions  the  equilibrium  is  very  far  in- 
deed toward  the  formation  of  the  hydrocarbonate  ion,  HCO3~, 
and  very  few  hydrogen  ions,  H+,  can  remain  in  the  solution  so 
long  as  any  carbonate  ions  are  present.  Since  many  more  hydro- 
gen ions  separate  from  carbonic  acid,  H2CO3,  when  that  is  present, 
than  separate  from  the  hydrocarbonate  ion,  HCO3~,  carbonate 


376  A  TEXTBOOK  OF  CHEMISTRY 

ions,  COs",  and  carbonic  acid,  H2CO3,  cannot  exist  to  any  large 
extent  in  the  same  solution.  It  will  be  seen  from  the  above 
that  carbonic  acid  cannot  be  formed  in  the  solution  in  suffi- 
cient amount  for  the  rapid  escape  of  carbon  dioxide  l  until  all  of 
the  carbonate  ions  present  have  been  converted  to  hydrocarbon- 
ate  ions.  In  other  words,  enough  hydrochloric  acid  must  be 
added  to  complete  the  reaction  : 

Na2CO3  +  HC1  =  NaHCO3  +  NaCl 

before  carbonic  acid  can  be  formed  in  sufficient  amount  for  the 
rapid  escape  of  carbon  dioxide.  The  addition  of  a  small  addi- 
tional amount  of  acid,  however,  causes  the  fourth  and  fifth 
reactions  to  occur  and  the  escape  of  carbon  dioxide  begins. 
From  this  point  on  the  gas  will  escape  in  almost  exact  proportion 
to  the  acid  added,  since  the  equilibrium  is  far  to  the  right  in  the 
first,  fourth  and  fifth  reactions. 

It  is  clear  from  the  above  discussion  that  the  formation  of 
a  volatile  product,  which  escapes  as  a  gaseous  phase,  has  the 
same  effect  on  the  course  of  a  reaction  when  it  is  a  secondary 
product,  formed  by  dissociation,  as  when  it  is  formed  directly 
by  the  ionization  reactions,  as  was  the  case  with  hydrochloric 
acid. 

Effect  of  Insolubility.  When  a  compound  separates  as  a  solid 
phase  from  a  solution,  the  effect  upon  the  equilibrium  is  exactly 
analogous  to  the  effect  of  volatilit^.  Thus  in  the  reactions  which 
occur  on  mixing  a  solution  of  salt  with  a  solution  of  silver  nitrate  : 

NaCl  ^±  Na+  +  Cl~ 
AgNO3  ^±  Ag+  +  NO3 
Ag+  +  Cl-  ^±  AgCl 


1  A  solution  of  sodium  bicarbonate,  NaHCO3,  contains  some  car- 
bonic acid,  H2CO3,  formed  by  the  reaction  : 

Na++  HCO3-  +  H+  +  OH-  =  H2CO3  +  Na++  OH~  . 

Such  a  solution  will  lose  carbon  dioxide  on  boiling  or  on  exposure 
to  the  air. 


SOLUBILITY  PRODUCT  377 

the  fact  that  silver  chloride  is  only  very  slightly  soluble  causes 
it  to  separate  from  the  solution  as  a  precipitate  and  shifts  the 
equilibrium  of  the  first  three  reactions  toward  the  formation 
of  this  compound.  The  removal  of  a  substance  from  the  mixture 
as  a  solid  phase  has  exactly  the  same  effect  upon  the  various 
equilibria  involved  as  the  removal  of  a  volatile  product. 

Effect  of  a  Common  Ion.  Solubility  Product.  When  a  sub- 
stance having  one  of  the  ions  of  a  difficultly  soluble  uni-uni- 
valent  salt,  that  is,  a  salt  such  as  silver  chloride,  AgCl,  in  which 
both  the  metallic  and  acid  radicals  are  univalent,  is  added  to  a 
saturated  solution  of  the  salt,  some  of  the  difficultly  soluble  salt 
will  usually  be  precipitated.  Thus  the  addition  of  a  few  drops 
of  a  solution  of  silver  nitrate  to  a  saturated  solution  of  silver 
chloride  will  cause  the  separation  of  some  silver  chloride  from 
the  solution  and  the  addition  of  hydrochloric  acid  will  also 
cause  precipitation  in  a  saturated  solution  of  silver  chloride. 

The  precipitation  depends  on  the  following  reactions : 


HC1  ^±  H+  +  Cl- 

In  the  first  reaction  the  quantities  of  the  three  substances, 
silver  chloride,  AgCl,  silver  ion,  Ag+,  and  chloride  ion,  Cl~,  must 
be  present  in  definite  amounts  in  a  saturated  solution  of  silver 
chloride  at  a  given  temperature.  At  18°  the  amount  of  silver 
chloride  in  the  three  forms  dissolved  by  a  liter  of  water  is  1.6 
milligrams  of  which  about  0.4  of  a  milligram  is  chlorine.  If, 
now,  hydrochloric  acid  is  added,  the  number  of  chloride  ions 
must  be  largely  increased,  and  this  will  cause  the  formation  of 
more  un-ipnized  silver  chloride.  As  the  solution  is  already 
saturated  with  that  compound,  the  silver  chloride  formed  will 
be  precipitated. 

Under  such  conditions  as  these,  when  a  solid,  difficultly 
soluble  uni-univalent  salt  is  in  equilibrium  with  a  solution  con- 
taining a  slight  excess  of  one  of  its  ions,  the  equation : 

CAg+  X*CC1-  =  a  constant 


378  A  TEXTBOOK  OF  CHEMISTRY 

has  been  found  to  be  true.  CAg+  arid  Ca-  are  the  concentrations 
of  the  ions  of  the  salt  in  the  solution.1 

If,  for  instance,  we  add  enough  hydrochloric  acid  so  that  the 
solution  is  y^Vs"  normal  for  hydrochloric  acid,  one  liter  will  con- 
tain 35.5  +  0.4  =  35.9  milligrams  of  chloride  ions,  Cl~ ,  or  nearly 
ten  times  as  much  as  before,  since  at  these  dilutions  both  the 
silver  chloride  and  hydrochloric  acid  will  be  almost  completely 
ionized.  Under  these  conditions,  in  order  that  the  product 
CAg+  X  CGI-  may  remain  constant  the  quantity  of  the  silver 
ions  must  become  -^  as  large  as  before.  Since  the  silver  chlo- 
ride is  almost  completely  ionized  at  these  concentrations,  it  will 
be  seen  that  the  addition  of  this  small  quantity  of  hydrochloric 
acid  will  reduce  the  quantity  of  silver  in  the  solution  to  about 
one  tenth  of  the  original  amount.  The  importance  of  these  re- 
lations for  the  precipitation  of  difficultly  soluble  salts  in  quantita- 
tive analysis  will  be  easily  seen.  This  principle  is  often  called 
the  constancy  of  the  solubility  product,  and  may  be  stated  as 
follows  :  In  any  dilute  aqueous  solution  saturated  with  a  slightly 
soluble  uni-univalent  salt  the  product  of  the  concentrations  of  the 
ions  of  this  salt  is  constant  at  a  given  temperature. 

In  some  cases  the  addition  of  a  salt  having  a  common  ion  will 
cause  precipitation  in  a  saturated  solution  containing  a  uni- 
bivalent  salt,  such  as  lead  chloride,  PbCl2,  or  in  one  containing 
a  bi-bivalent  salt,  such  as  calcium  oxalate,  CaC2O4,  but  there  are 
many  other  cases  in  which  such  precipitation  does  not  occur. 
No  general  rules  governing  the  conduct  of  such  salts  can  be  given 
in  the  present  state  of  our  knowledge.  In  some  cases  this  failure 
of  precipitation  may  be  due  to  the  formation  of  intermediate 
ions,  such  as  PbCl~.  In  others  it  is  probably  caused  by  the  for- 
mation of  more  or  less  stable  complex  ions. 

Formation  of  Complex  Ions.  If  solutions  of  silver  nitrate, 
AgNOs,  and  potassium  cyanide,  KCN,  are  mixed  in  equivalent 

1  This  relation  can  be  derived  theoretically  from  the  ionization 
reaction :  _,,  ^  . 

AgCl  ^  Ag+  +  Cl- 

if  it  is  assumed  that  in  very  dilute  solutions  the  osmotic  pressures 
of  the  ions  are  proportional  to  their  concentrations. 


COMPLEX  IONS.    IONIZATION  379 

proportions,  nearly  all  of  the  silver  ions  and  nearly  all  of  the 
cyanide  ions  will  be  removed  from  the  solution  as  a  precipitate, 
in  accordance  with  the  equation  : 

(AgN03  +  KCN  ^±  AgCN  +  KNO3 
Silver 
Cyanide 

If,  however,  more  of  the  potassium  cyanide  solution  is  added, 
the  silver  cyanide  will  dissolve.  This  seems  to  be  in  direct  con- 
tradiction to  the  principle  of  the  solubility  product  given  above. 
A  more  careful  examination  shows  that  the  apparent  contradic- 
tion is  due  to  the  fact  that  the  solution  no  longer  contains  an 
appreciable  number  of  silver  ions,  Ag+.  This  can  be  shown  in 
three  ways : 

1.  Sodium  chloride,  NaCl,  will  cause  no  precipitate  of  silver 
chloride,  AgCl,  to  form  in  the  solution. 

2.  If  an  electric  current  is  passed  through  the  solution,  the 
silver  in  the  solution  is  carried  toward  the  anode,  not  toward  the 
cathode.     This  shows  that  the  silver  atoms  form  part  of  complex 
ions  which  carry  negative  charges  instead  of  existing  in  the  solu- 
tion as  silver  ions.     If  the  silver  were  in  the  form  of  the  ion,  Ag+ , 
the  electric  current  would  carry  it  toward  the  cathode,  or  negative 
pole. 

3.  By  evaporating  the  solution  a  definite  compound,  potassium 
silver  cyanide,  KAgC2N2,  may  be  crystallized  from  it.     Trans- 
ference experiments  have  shown  that  the  ions  of  this  compound 
are  K+  and  AgC2N2~.     (See  p.  320.) 

Many  other  complex  ions  are  known  which  are  very  different 
in  their  properties  from  the  ions  which  unite  to  form  them. 

It  is  clear  from  this  illustration  that  the  principle  of  the  solu- 
bility product  depends  on  the  character  of  the  ions  which  are 
actually  present  and  not  on  the  amount  of  a  given  element  which 
may  be  present  in  the  solution. 

Degree  of  lonization.  If  two  electrodes  in  circuit  with  a 
battery  and  ammeter  are  placed  at  the  ends  of  a  narrow,  rectangu- 
lar cell  filled  with  distilled  water  (Fig.  93)  and  a  concentrated 
solution  of  potassium  chloride  is  put  in  the  bottom  of  the  cell, 


380 


A  TEXTBOOK  OF  CHEMISTRY 


the  ammeter  will  indicate  the  passage  of  an  electric  current. 

If,  now,  the  solution  is  stirred  so  that  the  potassium  chloride  is 

. +     uniformly   mixed   with 

I, r —  the   water,    it    will   be 

^^_J|        n/  y  seen  that   the   current 

increases.  Since  the 
amount  of  potassium 
chloride  in  the  cell  is 
not  changed  by  mixing 


Fig.  93 


the  solution  with  the  water  above  it,  it  is  evident  that  a  given 
amount  of  the  salt  is  more  effective  in  conveying  the  current  in 
a  dilute  solution  than  in  a  concentrated  one.  This  is  explained 
by  the  kinetic  theory  and  theory  of  ionization  by  supposing, 
first,  that  only  the  ions,  K+  and  Cl~,  take  part  in  the  conduc- 
tivity of  the  solution,  and  second,  that  in  the  ionization  reaction  : 


the  equilibrium  is  displaced  to  the  right  by  dilution  because  the 
potassium  and  chloride  ions  meet  each  other  to  unite  less  fre- 
quently in  the  dilute  solution,  while  the  tendency  to  separate 
into  ions  is  about  the  same  in  one  solution  as  in  the  other.  If 
the  conductivity  of  solutions  of  potassium  chloride  at  increasing 
dilution  is  measured  in  this  way,  the  results  given  in  the  follow- 
ing table  are  obtained.  Of  course,  for  the  dilute  solutions  it 
would  be  necessary  to  calculate  from  the  conductivity  of  a 
relatively  small  quantity  of  the  solution  what  the  conductivity 
of  the  whole  quantity  of  potassium  chloride  would  be. 

CONDUCTANCE  OF  SOLUTIONS  OF  POTASSIUM  CHLORIDE 

N  Potassium  Chloride  (74.5  grams  in          1  liter  )         98.28  mhos1  0.76 

N/10  Potassium  Chloride  (74.5  grams  in        10  liters)  111.97  mhos  0.865 

N/100  Potassium  Chloride  (74.5  grams  in       100  liters)  122.37  mhos  0.945 

N/1000  Potassium  Chloride  (74.5  grams  in    1000  liters)  127.27  mhos  0.983 

N/10000  Potassium  Chloride  (74.5  grams  in  10000  liters)  129.00  mhos  0.996 

N/oo  Potassium  Chloride  (74.5  grams  in      oo     liters)  129.5    mhos  1.000 


1  The  conductance  in  reciprocal  ohms  is  the  reciprocal  of  the  resist- 
ance in  ohms.     The  unit  for  conductance  is  one  mho. 


DEGREE   OF   IONIZATION  381 

Under  A  is  given  the  conductance  in  reciprocal  ohms  *  of  one 
gram  molecule  of  potassium  chloride  in  the  volumes  of  solution 
stated,  when  placed  between  two  electrodes  one  centimeter  apart 
and  sufficiently  large  to  contain  the  whole  solution  between 
them.  For  a  normal  solution  the  electrodes  might  be  25  X  40 
cm.,  giving  a  surface  of  1000  sq.  cm. 

It  will  be  seen  that  the  values  for  the  conductances  with  in- 
creasing dilution  form  a  converging  series  from  which  a  value  of 
129.5  for  infinite  dilution  can  be  calculated.  In  accordance  with 
the  theory  that  the  electric  current  is  carried  only  by  the  ions  and 
that  at  infinite  dilution  the  compound  is  completely  ionized, 
it  is  easy  to  calculate  the  degree  of  ionization  by  dividing  the 
conductance  for  any  given  concentration  by  the  conductance 
at  infinite  dilution.  Thus  the  fraction  of  a  normal  solution  of 

98.28 
potassium  chloride  in  the  form  of  ions  is       '       =  0.76.     The 


values  given  in  the  last  column  of  the  table  above  have  been 
calculated  in  this  manner. 

In  a  previous  chapter  it  was  pointed  out  (p.  112)  that  a  solu- 
tion containing  46  grams  of  alcohol  in  10  liters  of  water  freezes 
at  —  0.184.  One  containing  74.5  grams  of  potassium  chloride 
in  10  liters  freezes  at  —  0.343°.  It  would  seem  from  this  that  if 
the  potassium  chloride  were  completely  ionized,  the  depression 
of  the  freezing  point  would  be  twice  that  of  the  alcohol,  or  0.368°. 
This  would  correspond  to  an  increase  of  0.184°  in  the  depression 
of  the  freezing  point.  Comparing  this  with  the  increased  de- 


° 


pression  which   is   observed    (0.343°  —  0.184°)  we  have 

=  0.864  as  the  fraction  ionized.    This  is  in  very  close  agreement 
with  the  results  found  by  the  conductivity  method. 

Cane  sugar  is  hydrolyzed  by  dilute  acids  to  a  mixture  of  glu- 
cose and  fructose  (p.  334).  It  is  found  that  a  given  amount  of 
hydrochloric  or  nitric  acid  causes  a  much  more  rapid  hydrolysis 
than  an  equivalent  amount  of  acetic  acid.  If  we  assume  that 
the  rate  of  hydrolysis  is  proportional  to  the  number  of  hydrogen 
ions  present,  it  is  possible  to  calculate  from  series  of  experiments 


382  A  TEXTBOOK  OF  CHEMISTRY 

with  different  acids  the  relative  ionization  of  the  acids.  The 
results  of  such  experiments  are,  again,  in  general  agreement  with 
the  results  obtained  by  the  conductivity  and  freezing-point 
methods. 

The  three  methods  all  indicate  that  there  are  very  great  differ- 
ences in  the  degree  of  ionization  of  different  compounds.  The 
fact  that  the  results  obtained  by  three  methods  so  radically 
different  are  in  approximate  agreement  is  very  strong  evidence 
that  the  three  effects  have  a  common  basis,  and  no  satisfactory 
theory  other  than  that  of  ionization  has  been  proposed  to  account 
for  the  phenomena  observed.  The  lack  of  complete  agreement 
indicates  that  some  factors  which  are  not  yet  entirely  understood 
modify  the  effects,  just  as  the  mutual  attraction  of  the  molecules 
of  gases  prevent  them  from  obeying  exactly  the  laws  of  Avo- 
gadro  and  of  Boyle. 

The  following  table  gives  the  degrees  of  ionization  of  a  number 
of  common  substances.  In  the  measurement  and  calculation 
of  these  values  some  factors  which  have  not  been  discussed  are 
involved,  but  the  fundamental  principle  used  is  a  comparison  of 
the  conductivity  of  the  solution  for  which  the  degree  of  ionization 
is  given  with  the  conductivity  of  the  same  substance  when  it  is 
completely  ionized. 

FRACTIONS  IONIZED 

As  the  mobilities  of  the  H+and  OH~  ions  change  with  the  ion  con- 
centrations, the  conductance  ratio  (  A/  Aoo  )  does  not  correctly  represent 
the  fractions  ionized  for  strong  acids  and  bases.  Recent  investigations 
have  shown  that  these  substances  are  ionized  to  about  the  same  extent 
as  salts  of  the  same  type,  and  values  for  them  are  so  given  in  the  table. 
With  salts  of  acids  and  bases  having  bivalent  ions,  the  ionization  rela- 
tions are  complicated  by  the  presence  of  intermediate  and  complex  ions, 
for  instance  HgCl2  gives  besides  Hg++  and  Cl~,  the  ions  HgCl~  and 
HgCl4=.  The  figures  given  must  therefore  be  regarded  only  as  relative 
measures  of  the  tendencies  of  these  substances  to  form  ions. 


DEGREE   OF  IONIZATION 


383 


TENTH 

HUNDREDTH 

SUBSTANCE 

IONS 

FORMULA 
SOLUTIONS 

FORMULA 
SOLUTIONS 

V  =  10 

V  =  100 

Nitric  Acid      .... 

H+,  N03- 

0.86 

0.94 

Hydrochloric  Acid   . 

H+,  Cl- 

0.86 

0.94 

Hydrobromic  Acid  .     . 

H  +,  Br~ 

0.86 

0.94 

Hydroiodic  Acid       .     . 

H+,  I- 

0.86 

0.94 

Chloric  Acid   .... 

H+,  C103- 

0.86 

0.94 

Perchloric  Acid   .     .     . 

H+,  C1O4- 

0.86 

0.94 

Permanganic  Acid    .     . 

H+,  MnO4- 

0.86 

0.94 

Tartaric  Acid 

H+,  HC4H406- 

0.098 

0.31 

Acetic  Acid     .... 

H+,  C2H302- 

0.013 

0.043 

Hydrocyanic  Acid    .     . 

H+,  CN- 

0.00008 

0.00026 

Boric  Acid      .... 

H+,  H2B03- 

0.00008 

0.00026 

Hydrosulfuric  Acid  .     . 

H+,  HS- 

0.00095 

0.00031 

Sulfurous  Acid     .     .     . 

H+,  HS03- 

0.50 

0.70 

Carbonic  Acid 

H+,  HCO-r 

0.0017 

0.0055 

Phosphoric  Acid       .     . 

H+,  H2PO4~ 

0.28 

0.64 

Phenol        .     .     . 

H+  ,  C6H5O- 

0.00011 

0.00036 

Sulfuric  Acid  .... 

i  (H+,  H+,  S04=) 

0.608 

0.832 

Oxalic  Acid     .... 

i  (H+,  H+,  C204=) 

0.17 

0.398 

Sodium  Hydroxide  .     . 

Na+,  OH- 

0.86 

0.94 

Potassium  Hydroxide  . 

K+,  OH- 

0.86 

0.94 

Tetra  methyl       Ammo- 

nium Hydroxide  .     . 

N(CH3)4+,  OH- 

0.85 

0.94 

Barium  Hydroxide  . 

!(Ba++,OH-,OH-) 

0.76 

0.88 

Ammonium  Hydroxide1 

NH4+,  OH- 

0.013 

0.042 

Water  2 

H+,  OH~ 

0.0000001 

0.0000001 

1  This  gives  the  fraction  of  the  total  ammonia  in  the  solution 
which  is  in  the  form  of  ammonium,  NH4+,  and  hydroxide,  OH~, 
ions.    Actually  much  of  the  ammonia  is  present  as  NH3  and  the 
proportion  of  real  NH4OH  ionized  is  much  larger  than  that  given 
in  the  table. 

2  The  value  for  water  gives  the  fraction  of  a  mol  of  water  in 
one  liter,  which  is  in  the  form  of  ions  at  25°.     For  comparison  with 
the  other  values  in  the  table  the  values  under  V  —  10  must  be 
divided  by  10  and  those  under  V  —  100  must  be  divided  by  100. 
Thus  one  liter  of  tenth-normal  hydrocyanic  acid  contains  0.000008 
mol  in  the  form  of  ions  and  one  liter  of  hundredth-normal  acid 
contains  0.0000026  mol  in  that  form,  while  one  liter  of  pure  water 
contains  0.0000001  mol. 


384 


A  TEXTBOOK  OF  CHEMISTRY 


FRACTIONS  IONIZED  —  Continued 


SUBSTANCE 

IONS 

TENTH 
FORMULA 
SOLUTIONS 
V  =  10 

HUNDREDTH 
FORMULA 
SOLUTIONS 
V  =  100 

Sodium  Chloride 

Na+,  Cl- 

0.852 

0.935 

Potassium  Chloride     ... 

K+,  Cl- 

0.855 

0.941 

Ammonium  Chloride    . 

NH4+,  Cl- 

0.852 

0.936 

Sodium  Nitrate  .     .     . 

Na+,  NO3~ 

0.832 

0.932 

Potassium  Nitrate  . 

K+,  NO3- 

0.824 

0.935 

Silver  Nitrate      . 

Ag+,  N03- 

0.816 

0.931 

Potassium  Chlorate 

K+,  C1O3- 

0.824 

0.933 

Sodium  Acetate       .     . 

Na+,  C2H3O2~ 

0.795 

0.914 

Potassium  Cyanide 

K+,  CN- 

0.84 

0.93 

Sodium  Bicarbonate     . 

Na+,  HCO3~ 

0.84 

0.93 

Potassium  Sulfate    .     . 

i  (K+,  K+,  S04=) 

0.724 

0.870 

Sodium  Sulfate    .     .     . 

i(Na+,Na+,SO4=) 

0.704 

0.857 

Normal    Sodium    Car- 

i  (Na+,  Na+,  CO3=) 

0.71 

0.86 

Calcium  Sulfate  .     .    . 

i  (Ca++,  S04=) 

0.625 

Zinc  Chloride      .     .     . 

i  (Zn++,Cl-,Cl-) 

0.71 

0.86 

Zinc  Sulfate    .     .    .    . 

i  (Zn++,  S04=) 

0.405 

0.633 

Copper  Sulfate    .    .    . 

}  (Cu++,  S04=) 

0.396 

0.629 

Mercuric  Chloride   .     . 

J(Hg++,Cl-,Cl-) 

0.01 

0.03 

(This  table  was  prepared  by  Dr.  D.  A.  Maclnnes.) 

Effect  of  Degree  of  lonization.  Neutralization.  When  a 
volatile  product  escapes  from  a  mixture,  or  when  a  solid  substance 
is  precipitated,  the  equilibria  of  the  reactions  which  lead  to  the 
formation  of  such  compounds  are  shifted  in  such  a  way  as  to 
promote  their  formation.  From  the  table  which  has  just  been 
given  it  is  evident  that  certain  ions  cannot  exist  in  any  number  in 
solutions  which  contain  certain  other  ions.  There  can  be  very 
few  hydrogen  ions,  H+,  in  solutions  containing  hydroxide,  OH~, 
hydrosulfide,  HS~,  or  hydrocarbonate,  HCOs",  ions,  or  in  the 
presence  of  the  ions  of  any  of  the  weak  acids.  If  the  solutions  of 
two  substances  giving  hydrogen  and  hydroxide  ions  are  mixed 
in  equivalent  amounts,  the  two  ions  unite  to  form  water  and 


HYDROLYSIS  385 

the  resulting  solution  is  neutral,  if  both  of  the  compounds  have 
the  same  degree  of  ionization  : 

HCI  ;£  H+  +  cr 

NaOH  ^±  Na+  +  OH~ 
H+  +  OH-  ^±  H2O 

The  equilibrium  of  the  last  reaction  is  so  far  toward  the  forma- 
tion of  water  that  both  hydrogen  and  hydroxyl  ions  and  also 
practically  all  of  the  hydrochloric  acid  and  sodium  hydroxide 
disappear  from  the  solution.  This  is  the  ordinary  reaction  of 
neutralization.  A  solution  is  neutral  when  the  number  of  hy- 
drogen, H+,  and  hydroxide,  OH~,  ions  is  equal. 

An  illustration  of  the  effect  of  adding  hydrochloric  acid  to  a 
solution  containing  carbonate  ions  was  given  above  (p.  375). 

Hydrolysis.  It  will  be  seen  from  the  table  that  the  ionization 
of  carbonic  acid,  ^COs,  to  hydrogen,  H+,  and  hydrocarbonate, 
HCOs",  ions  is  very  slight.  The  further  ionization  of  hydrocar- 
bonate ions  to  hydrogen,  H+,  and  carbonate,  CO3=,  ions  must 
be  almost  vanishingly  small  in  the  presence  of  an  excess  of  hy- 
drogen ions.  A  solution  of  sodium  carbonate,  Na2CO3,  which 
may  ionize,  at  first,  as  follows : 

Na2CO3  ^±  Na+  +  Na+  +  CO3= 
immediately  gives  with  the  ions  of  water 

COr  +  H+  +  OH-  ^±  HCOr  +  OH- 

This  formation  of  hydrocarbonate  ions,  HCO*-,  reduces  the 
number  of  hydrogen  ions  and  the  solution  must  react  alkaline 
because  of  excess  of  hydroxide  ions,  OH",  present.  Similar  re- 
actions, due  to  the  ions  of  the  water  and  called  hydrolysis, 
occur  with  the  alkali  metal  salts  of  all  of  the  very  weak  acids 
and  especially  with  the  salts  of  weak  dibasic  and  tribasic  acids, 
such  as  sulfides,  borates  and  phosphates. 

The  hydroxides  of  ferric  iron,  Fe(OH)3,  aluminium,  A1(OH)3, 
chromium,  Cr(OH)s,  and  of  many  other  elements  are  so  insoluble 
that  it  is  impossible  to  determine  the  degree  of  their  ionization, 


386  A  TEXTBOOK  OF  CHEMISTRY 

and  there  is  reason  to  think  that  the  ionization  of  the  hydroxyl 
in  basic  compounds  such  as  FeCl2OH,  which  probably  exist  in  so- 
lution, is  very  trifling.  In  a  solution  of  ferric  chloride,  therefore, 
such  reactions  as  the  following  occur  : 


FeCl2+  -f  OH"  +  H+  =  FeCl2OH  +  H+ 

Basic 
Ferric  Chloride 

The  resulting  solution,  which  contains  an  excess  of  hydrogen 
ions,  will  react  acid. 

These  facts  may  also  be  stated  in  the  form  :  Weak  acids  cannot 
neutralize  strong  bases  and  weak  bases  cannot  neutralize  strong 
acids  completely,  when  mixed  in  equivalent  proportions,  because 
of  the  hydrolysis  of  the  salts  formed.  It  is  chiefly  because  of 
these  relations  that  acids  and  bases  which  are  but  slightly  ionized 
in  their  solutions  are  called  "  weak."  The  degree  of  ionization 
furnishes  the  only  satisfactory  basis  for  classifying  acids  or 
bases  as  "  weak  "  or  "  strong." 

Illustration  of  the  Strength  of  Acids.  If  solutions  of  potassium 
iodide  (KI,  6  grams  per  liter)  and  potassium  bromate  (KBrO3, 
1  gram  per  liter)  are  mixed  in  equal  proportions,  no  reaction 
occurs  between  them,  but  on  the  addition  of  an  acid,  iodine  is 
liberated  and  colors  the  solution  yellow  or  brown  : 

6  HI  +  HBrO3  =  HBr  +  3  I2  +  3  H2O 

If  the  acid  is  dilute,  the  reaction  is  sufficiently  slow  so  that  its 
progress  can  be  noted  by  a  slow  change  in  the  color.  The  rate 
of  the  change  is  proportional  to  the  number  of  hydrogen  ions 
present.  If  equal  volumes  of  the  iodide-bromate  solution  are 
placed  in  three  glasses  and  there  is  added  to  these,  respectively, 
5  cc.  of  tenth  normal  hydrochloric  acid  (3.65  g.  HC1  per  liter), 
5  cc.  of  tenth  normal  oxalic  acid  (5.3  g.  H2C2O4.2  H2O  per 
liter)  and  5  cc.  of  tenth-normal  acetic  acid  (6.0  g.  HC2H3O2  per 
liter),  a  very  rapid  change  in  color  will  occur  in  the  first  solution, 
a  much  slower  change  in  the  second  and  a  very  slow  change, 
indeed,  in  the  third.  On  the  other  hand,  if  5  cc.  of  each  acid 
are  placed  in  three  beakers  and  5  cc.  of  tenth-normal  sodium  hy- 


USE   OF  INDICATORS  387 

droxide  are  added  to  each,  the  three  solutions,  containing  sodium 
chloride,  NaCl,  sodium  oxalate,  Na2C2O4,  and  sodium  acetate, 
NaC2H3O2,  respectively,  will  be  neutral  toward  phenol  phthalein 
or  litmus.  The  experiment  shows  that  while  the  three  acids 
differ  very  greatly  in  "  strength  "  as  shown  by  a  reaction  which 
depends  on  the  number  of  hydrogen  ions,  H+,  they  give  at  a 
given  dilution,  they  are  nearly  equal  in  their  power  of  neutral- 
ization, which  depends  for  these  acids  and  the  indicator  chosen 
on  the  number  of  hydrogen  ions  which  can  be  obtained  from 
them  by  complete  ionization.  It  should  be  remembered,  how- 
ever, that  sodium  oxalate  and  sodium  acetate  are  not,  strictly 
speaking  neutral  salts.  The  reason  for  this  lack  of  neutrality 
will  be  clearer  after  a  study  of  the  following  paragraph. 

Use  of  Indicators.  It  has  been  stated  (p.  122)  that  indicators 
are  colored  substances  which  exist  in  two  forms,  one  of  which 
is  stable  in  the  presence  of  hydrogen,  H+,  ions  while  the  other 
is  stable  in  the  presence  of  hydroxide,  OH~~,  ions.  A  neutral 
solution  is  one  in  which  the  numbers  of  hydrogen  and  hydroxide 
ions  are  equal.  It  has  been  pointed  out  in  the  table  on  p.  383 
that  the  number  of  mols  of  water  ionized  in  one  liter  is  0.0000001, 
or  10~7  at  25°.  Since  in  the  ionization  reaction 


we  must  have 

CH+  X  COH~  =  constant 

so  long  as  the  solution  consists  mostly  of  water,  it  follows  that 
in  dilute  solutions 

CH+  X  COH-  =  10~7  X  JO'7  =  10~14 

If  an  acid  is  added  to  water  in  sufficient  amount  to  increase 
the  concentration  of  the  hydrogen  ions  from  10~7  to  10~6,  the 
concentration  of  the  hydroxide  ions  must  fall  to  10~8.  Only 
0.0000009  gram-mol  of  hydrogen  ions  would  be  required  to 
produce  such  a  change  in  a  liter  of  water,  and  this  would  be 
given  by  about  0.001  cc.  of  tenth-normal  hydrochloric  acid, 
while  0.01  cc.  would  increase  the  hydrogen  ion  concentration 
to  ID'*. 


388 


A  TEXTBOOK  OF  CHEMISTRY 


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j 

USE   OF  INDICATORS  389 

Indicators  are  weak  acids  or  bases,  and  the  change  in  color 
does  not  usually  occur  in  an  exactly  neutral  solution.  The  table 
on  p.  388  gives  the  concentrations  of  hydrogen  and  hydroxide 
ions  at  which  the  change  in  color  occurs  for  some  of  the  more 
common  indicators. 

It  would  seem,  at  first  thought,  that  only  an  indicator  which 
changes  exactly  at  the  neutral  point  would  be  suitable,  but  this 
is  by  no  means  true.  In  titrating  a  strong  acid  with  a  strong 
base  a  few  hundredths  of  a  cubic  centimeter  of  tenth-normal 
acid  or  alkali  will  carry  the  concentration  of  hydrogen  or  hy- 
droxide ions  so  far  to  one  side  of  the  neutral  point  that  any  of 
the  indicators  for  which  the  acidity  is  between  10~6  and  10~9 
will  give  a  sharp  end  point. 

In  titrating  a  weak  acid,  such  as  acetic  acid,  HC2H3O2, 
with  a  strong  base,  such  as  potassium  hydroxide,  KOH,  as  the 
neutral  point  is  approached  the  potassium  acetate,  KC2HaO2, 
formed  is  much  more  highly  ionized  than  the  acetic  acid  and 
the  acetate  ions  carry  the  ionization : 

HC2H3O2±£H+  +  C2H3O2- 

far  to  the  left.  The  number  of  hydrogen  ions  then  becomes 
very  small,  while  a  considerable  amount  of  acetic  acid  is  still 
unneutralized.  Under  such  conditions  the  change  in  color  for 
methyl  orange  or  methyl  red  may  appear  before  the  acid  is 
completely  neutralized  and  the  end  reaction  will  not  be  sharp, 
i.e.  the  change  in  color  will  appear  gradually  during  the  addition, 
sometimes,  of  a  cubic  centimeter  or  more  of  the  alkali.  But  if 
phenol  phthalein  is  used,  the  change  in  color  will  not  occur  till 
the  true  neutral  point  is  passed  and  then  a  very  slight  excess 
of  alkali  will  carry  the  concentration  of  the  hydroxide  ions  far 
beyond  the  neutral  point.  The  end  will  be  sharp  and  will  cor- 
respond closely  to  the  exact  neutralization  of  the  acid. 

When  a  weak  base  such  as  ammonium  hydroxide,  NH4OH, 
is  titrated  with  a  strong  acid,  as  hydrochloric  acid,  HC1,  the 
conditions  are  reversed  and  such  indicators  as  methyl  red, 


390  A  TEXTBOOK  OF  CHEMISTRY 

methyl  orange  or  cochineal,  which  change  color  in  a  faintly 
acid  solution,  are  most  suitable. 

With  very  weak  acids  or  bases  the  hydrolysis  of  the  salts 
formed  may  carry  the  acidity  or  alkalinity  of  the  normal  salt 
far  to  one  side  of  the  true  neutral  point.  An  illustration  of  the 
use  of  indicators  in  such  a  case  will  be  given  later  (p.  464). 

Systematic  Study  of  the  Metals.  In  discussing  the  metals  it 
is  natural  to  consider  first,  as  with  the  non-metals,  their  prepa- 
ration and  properties.  After  this  their  compounds  may  be  dis- 
cussed in  the  same  order  which  has  been  followed  in  the  consid- 
eration of  the  nonmetallic  elements :  oxides  and  hydroxides ; 
chlorides,  hypochlorites,  chlorates,  bromides,  iodides,  fluorides ; 
sulfides,  sulfites,  sulfates ;  nitrites,  nitrates,  phosphates,  arsen- 
ates  ;  carbonates,  bicarbonates,  salts  of  organic  acids,  cyanides  ; 
silicates ;  borates.  Most  of  the  metals  form  nearly  all  of  these 
classes  of  compounds,  but  only  those  compounds  which  are  of 
some  particular  scientific,  historical  or  industrial  importance  can 
be  mentioned  in  this  book. 

Before  taking  up  the  individual  metals  it  seems  desirable  to 
give  some  general  statements  with  regard  to  the  preparation  or 
metallurgy  of  the  metals  and  about  the  preparation  and  proper- 
ties of  the  various  classes  of  compounds. 

Metallurgy.  The  first  metals  used  by  man  were  those  which 
are  found  free  in  nature,  such  as  copper,  silver  and  gold.  The 
use  of  these  metals  and  their  alloys  marks  the  beginning  of  the 
"  Bronze  Age,"  which  reaches  back  into  prehistoric  times  — 
very  late  and  recent  when  we  consider  the  long  history  of  the 
race,  but  early  in  a  period  when  mankind  became  organized 
in  society  and  differentiated  sharply  from  the  animal  world. 
The  discovery  of  methods  of  reducing  iron  from  its  ores  by  means 
of  wind  furnaces  probably  occurred  during  the  period  of  history 
which  has  been  recorded  in  inscriptions,  but  still  so  early  that 
no  definite  record  is  to  be  found.  Inscriptions  in  Egypt  show 
that  iron  was  made  there  at  least  3000  years  ago.  From  that 
period  till  the  middle  of  the  nineteenth  century  all  methods  for 
industrial  metallurgy  depended  on  the  use  of  charcoal,  coal  or 


METALLURGY  391 

other  fuels  containing  carbon.  The  most  important  of  these 
methods  depended  on  the  direct  reduction  of  oxides  of  the  metals 
by  means  of  these  fuels,  while  a  few  depended  on  the  roasting 
of  a  sulfide  (mercury,  p.  485)  or  on  the  roasting  of  a  sulfide 
followed  by  reduction  through  the  interaction  of  an  oxide  with  a 
sulfide  or  sulfate  (copper,  p.  427,  and  lead,  p.  513). 

The  first  preparation  of  aluminium  by  Deville  in  1854  in 
sufficient  amount  to  demonstrate  its  valuable  properties  led  to  a 
strong  desire  to  secure  cheap  sodium  for  use  in  its  production. 
This  resulted  in  the  development  of  the  reduction  of  sodium 
carbonate  to  metallic  sodium  by  means  of  carbon  and  the  use 
of  the  latter  for  the  preparation  of  aluminium  from  its  chloride, 
Aids.  But  this  method  of  preparing  aluminium  did  not  attain 
any  considerable  industrial  importance. 

Electrolytic  methods  were  used  for  the  deposition  of  thin 
films  of  copper  and  other  metals  as  early  as  1836,  but  electrical 
methods  for  preparing  and  refining  metals  could  not  be  used 
on  a  large  scale  till  the  development  of  dynamos  during  the 
last  quarter  of  the  nineteenth  century  made  the  production  of 
relatively  cheap  electrical  energy  possible. 

The  first  industrial  use  of  an  electric  furnace  seems  to  have 
been  its  application  to  the  manufacture  of  aluminium  bronze  by 
the  Cowles  Brothers  of  Cleveland,  Ohio,  in  1884.  Their  labora- 
tory experiments  with  the  method  began  in  1882.  The  Hall 
method  for  the  electrolysis  of  aluminium  oxide  dissolved  in 
cryolite  soon  rendered  the  Cowles  furnace  industrially  worth- 
less for  this  particular  purpose,  but  electric  furnaces  are  now 
extensively  used  for  metallurgical  processes  and  for  many  other 
forms  of  chemical  manufacture. 

The  relatively  cheap  electrolytic  manufacture  of  aluminium 
has  not  only  given  an  abundant  supply  for  use  in  the  metallic 
form,  but  has  led  to  the  development  of  processes  for  the  pro- 
duction of  chromium  and  other  metals  by  heating  their  oxides 
with  aluminium. 

It  will  be  seen  from  this  brief  sketch  of  the  historical  develop- 
ment of  metallurgy  that  many  important  processes  are  of  very 


392  A  TEXTBOOK  OF  CHEMISTRY 

recent  origin.  It  is  only  a  very  few  years  since  several  metals 
which  now  have  very  important  industrial  uses,  were  scarcely 
more  than  scientific  curiosities.  Further  rapid  development 
along  such  lines  is  to  be  confidently  expected. 

Oxides.  All  metals,  without  exception,  may  be  combined 
with  oxygen  and  nearly  all  metals  combine  with  the  element 
directly  at  suitable  temperatures.  Even  the  so-called  "  noble  " 
metals,  gold,  silver  and  platinum  combine  with  oxygen  under 
some  conditions,  but  the  oxides  are  very  unstable,  are  easily 
decomposed  by  heat  alone  and  are  reduced  by  hydrogen  at 
ordinary  temperatures. 

A  few  metals,  especially  sodium,  potassium  and  barium,  com- 
bine with  oxygen  to  form  peroxides,  in  which  two  oxygen  atoms 


are  united,  as  in  Ba<Q  | ,  but  in  nearly  all  of  the  metallic  oxides 

Nj 

the  oxygen  seems  to  be  united  only  with  the  metal,  and  the 
valence  of  the  metal  is  apparent  from  the  formula  of  the  oxide. 
Nearly  all  nitrates,  carbonates  and  hydroxides  are  decom- 
posed by  heat  with  the  formation  of  oxides : 

2Pb(NO3)2  =  2PbO  +  4NO2  +  O2 

CaCO3  =  CaO  +  CO2 
Cu(OH)2  =  CuO  +  H2O 

It  is  doubtful  if  any  oxide  dissolves  appreciably  in  water  as 
an  oxide,  and  those  oxides  which  do  not  combine  with  water 
to  form  hydroxides  are  practically  insoluble. 

Hydroxides.  The  alkali  metals,  sodium,  potassium,  etc., 
and  the  alkali  earth  metals,  calcium,  strontium  and  barium 
decompose  water  with  the  formation  of  hydroxides  at  ordinary 
temperatures : 

Na  +  HOH    =  NaOH  +  H 
Ca  +  2HOH  =  Ca(OH)2  +  2H 

Magnesium  decomposes  water  at  100°.  Other  metals,  as  zinc 
and  iron,  which  decompose  water  at  higher  temperatures,  form 
oxides  instead  of  hydroxides,  though  it  is,  of  course,  possible 


SOLUBILITY  OF  SALTS  393 


that  the  latter  are  formed  at  first  and  immediately  decom- 
posed. 

Iron  is  converted  by  the  combined  action  of  water  and  air  into 
iron  rust,  a  combined  oxide  and  hydroxide  having  the  composi- 
tion of  the  mineral  limonite,  Fe2O3.2Fe(OH)3. 

Practically  all  hydroxides  of  the  metals  except  those  of  the 
alkali  and  alkali-earth  metals  are  insoluble  in  water.  For  this 
reason  the  hydroxides  of  nearly  all  other  metals  are  precipitated 
from  solutions  of  their  salts  by  solutions  of  sodium  or  potassium 
hydroxide.  In  a  few  cases,  especially  those  of  silver,  cuprous 
copper,  mercurous  and  mercuric  mercury,  these  hydroxides 
precipitate  an  oxide  instead  of  the  hydroxide,  doubtless  because 
the  hydroxides  of  these  metals  are  unstable : 

AgNO3  +  NaOH  =  [AgOH]  +  NaNO3 
[2AgOH]=Ag20  +  H20 

Solubility  of  Salts.  Practically  all  salts  of  the  alkali  metals 
(lithium,  sodium,  potassium,  ammonium)  are  soluble  in  water, 
the  only  important  exceptions  being  sodium  pyroantimonate, 
Na2H2Sb2O7.6  H2O,  potassium  and  ammonium  chloroplatinates, 
K2PtCl6,  and  (NH4)2PtCl6,1  potassium  perchlorate,  KC1O4,  and 
potassium  cobaltini trite,  K3Co(NO2)6,  or  potassium  silver 
cobaltinitrite,  K2AgCo(NO2)e.  Some  of  these  are,  however, 
more  soluble  than  those  salts  of  other  metals  which  are  usually 
counted  as  insoluble.  There  is  probably  no  salt  which  is 
wholly  insoluble  in  water. 

Nearly  all  salts  of  the  strong  monobasic  and  bibasic  acids 
are  also  soluble.  This  includes  chlorides,  bromides  and  iodides, 
fluorides,  chlorates  and  perchlorates,  sulfites  and  sulfates,  ni- 
trites and  nitrates.  The  most  important  exceptions  are  the 
chlorides,  bromides  and  iodides  of  silver,  cuprous  copper,  mer- 
curous mercury  and  lead,  AgCl,  AgBr,  Agl,  Cu2Cl2,  Cu2I2, 
Hg2Cl2,  Hg2I2,  PbCl2  (slightly  soluble),  PbBr2,  and  PbI2,  mer- 

1  Rubidium  and  caBsium  chloroplatinates,  Rb2PtCl«  and 
Cs2PtCl«  are  still  less  insoluble. 


394  A  TEXTBOOK  OF  CHEMISTRY 

curie  iodide,  HgI2,  calcium  fluoride,  CaF2,  barium  sulfite,  BaSO3, 
and  calcium,  strontium,  barium  and  radium  sulfates,  CaSO4 
(slightly  soluble),  SrSO4,  BaSO4,  RaSO4. 

Normal  salts  of  phosphoric,  H3PO4,  arsenious,  H3AsO3,  arsenic, 
H3AsO4,  carbonic,  H2CO3,  silicic,  H2SiO3,  etc.,  and  boric,  H3BO3, 
acids,  with  the  exception  of  those  of  the  alkalies,  are  insoluble. 
Sulfides  other  than  those  of  the  alkalies  are  either  insoluble  in 
water  or  are  hydrolyzed  with  the  formation  of  a  hydrosulfide 
(as  Ca(SH)2),  and  a  hydroxide,  or  of  hydrogen  sulfide,  H2S,  and 
an  insoluble  hydroxide,  such  as  A1(OH)3. 


CHAPTER  XXIII 
ALKALI  METALS :  LITHIUM,  SODIUM 

General  Properties  of  the  Alkali  Metals.  The  alkali  metals 
are  univalent  elements  which  combine  with  hydroxyl  to  form 
the  strongest  bases,  hydroxides  which  are  easily  soluble  in  water 
and  which  are  largely  ionized  in  solutions  of  moderate  concen- 
trations. They  are  the  most  active  of  the  metallic  elements,, 
decomposing  water  rapidly  at  ordinary  temperatures  and  tar- 
nishing almost  instantly  in  ordinary  air,  owing  to  the  formation 
of  a  film  of  hydroxide. 

Their  affinity  for  the  halogens  is  also  so  strong  that  sodium 
and  potassium  have  often  been  used  to  decompose  halides  for 
the  preparation  of  metals  and  other  elements. 

As  has  been  stated  in  the  last  chapter,  nearly  all  salts  of  the 
alkali  metals  are  soluble  in  water.  Normal  salts  of  weak  acids, 
such  as  the  sulfides,  carbonates,  cyanides,  phosphates,  silicates 
and  borates  are  hydrolyzed  by  water,  and  their  solutions  have  a 
strongly  alkaline  reaction. 

The  metals  of  the  group  have  a  low  specific  gravity,  lithium, 
sodium  and  potassium  being  lighter  than  water.  Their  melting 
points  range  from  186°  for  lithium  to  26.5°  for  caesium  and  the 
boiling  points,  from  above  a  red  heat  for  lithium  and  742°  for 
sodium  to  270°  for  caesium. 

Lithium,  Li,  6.94,  is  usually  considered  as  one  of  the  rarer 
elements  and  of  comparatively  little  importance.  It  is  found 
in  a  number  of  silicates  and  in  small  amount  in  practically  all 
natural  waters.  The  metal  has  a  specific  gravity  of  only  0.51. 
Hydrogen  and  helium  are  the  only  elements  which  are  lighter 
than  lithium  when  in  the  solid  or  liquid  state.  The  metal  com- 
bines with  hydrogen  to  form  the  hydride,  LiH,  and  with  nitrogen 

395 


396  A  TEXTBOOK  OF  CHEMISTRY 

to  form  the  nitride,  Li3N.  It  has  been  used  in  the  separation  of 
argon  from  the  .atmosphere  because  of  its  strong  affinity  for 
nitrogen. 

Lithium  carbonate,  I^COs,  may  be  decomposed  to  lithium 
oxide,  Li2O,  and  carbon  dioxide  by  heating  it  in  a  current  of 
hydrogen.  Both  the  carbonate  and  the  phosphate,  LisPO-i,  are 
difficultly  soluble  in  water.  In  these  properties  lithium  re- 
sembles magnesium,  the  second  element  of  the  second  group, 
rather  than  the  other  alkali  metals.  Beryllium,  the  first  ele- 
ment of  the  second  group,  approaches  aluminium  in  its  proper- 
ties in  a  similar  manner. 

Lithium  Urate,  LiC5H3O3N4,  is  soluble  in  water,  and  this  fact 
led  physicians  to  the  belief  that  the  administration  of  lithium 
carbonate  or  of  natural  waters  containing  lithium  would  be 
beneficial  to  patients  suffering  from  rheumatism  or  gout  and 
they  have  been  much  employed  as  remedies  in  those  diseases. 
A  more  careful  study  has  shown  that  these  compounds  are  worth- 
less for  such  a  purpose,  but  the  ingestion  of  large  quantities  of 
water  with  the  lithium  compounds  probably  exerts  a  beneficial 
effect. 

Lithium  compounds  impart  to  the  Bunsen  flame  a  brilliant 
red  color  and  give  a  spectrum  of  two  red  lines,  one  of  which  is 
very  bright. 

Atomic  Weight  of  Lithium.  Law  of  Dulong  and  Petit.  It 
has  been  pointed  out  (p.  92)  that  the  most  satisfactory  method 
of  selecting  the  true  atomic  weight  of  an  element  consists  in 
finding  the  weight  of  the  element  contained  in  a  gram-molecular 
volume  (22.4  liters  at  0°  and  760  mm.)  of  that  gaseous  compound 
which  contains  the  smallest  quantity  of  the  element  in  this 
volume.  But  lithium  forms  no  compound  whose  weight  in  the 
gaseous  form  has  been  determined,  and  a  considerable  number 
of  other  elements  form  no  compounds  which  can  be  converted 
into  gases  without  decomposition.  The  atomic  weights  of  such 
elements  must,  of  course,  be  selected  in  a  different  manner. 
For  this  purpose  the  law  of  Dulong  and  Petit,  discovered  in 
1819,  has  been  useful.  These  chemists  found  that  the  quantity 


LAW  OF  DULONG  AND  PETIT 


397 


of  heat  required  to  raise  the  temperature  of  one  gram-atom  of 
an  element  one  degree  is  approximately  6.6  calories.  If  this 
quantity  of  heat  is  applied  to  7  grams  of  lithium  or  to  65  grams 
of  zinc  or  to  200  grams  of  mercury,  it  will,  in  each  case,  raise  the 
temperature  one  degree. 

The  law  is  also  frequently  stated  that  the  specific  heat  of  an 
element  multiplied  by  its  atomic  weight  is  a  constant  quantity. 
The  following  table  will  make  this  clear : 


ELEMENT 

SPECIFIC 
HEAT 

ATOMIC 
WEIGHT 

SP.  HT.  X  AT.  WT. 

Lithium                 

0.94 

7. 

66 

Graphite  (at  11°)      .     . 
Graphite  (at  977°)    .... 
Silicon 

0.16 

0.467 
0  16 

12. 
12. 

28.4 

1.9 
5.6 
4.5 

Calcium      
Zinc 

0.17  . 
0093 

40. 
654 

6.8 
6.1 

0.084 

80. 

6.7 

0.033 

200. 

6.7 

0.03 

207. 

6.2 

It  will  be  seen  from  the  table  that  graphite  and  silicon  depart 
rather  widely  from  the  law,  though  the  former  approaches  it 
more  closely  at  high  temperatures.  All  of  the  metallic  elements 
and  all  elements  having  atomic  weights  above  40  conform  ap- 
proximately to  the  law.  The  law  is  at  best,  however,  only 
approximate  and  is  of  service  only  in  selecting  between  rather 
widely  divergent  possible  values  for  an  atomic  weight.  Thus 
the  atomic  weight  of  calcium  might  be  20,  40  or  60,  according 
as  the  formula  of  the  chloride  is  CaCl,  CaCl2  or  CaCl3.  But  of 
these  three  values  only  an  atomic  weight  of  40  agrees  with  the 
law. 

The  laws  of  Avogadro  and  of  Dulong  and  Petit  have  usually 
been  considered  as  independent  and  wholly  unrelated.1  A 


1  See,  however,  G.  N.    Lewis,  J.  Am.  Chem.  Soc.  29,  1165  and 
1516  (1907). 


398  A  TEXTBOOK  OF   CHEMISTRY 

little  consideration,  however,  shows  us  that  if  we  accept  the 
kinetic-molecular  theory,  this  is  not  the  case.  At  foundation 
Avogadro's  law  depends  on  the  fact  that  molecules  of  different 
weights  exchange  energies,  when  in  collision  with  each  other 
or  with  the  walls  of  the  containing  vessel  at  a  given  tempera- 
ture, in  such  a  manner  that  the  average  value  of  J  mv2  (m  = 
mass,  v  =  velocity)  is  constant  and  is  independent  of  the 
weight  of  the  molecule.  The  law  of  Dulong  and  Petit  must 
depend  on  a  similar  property  of  the  atoms  of  the  elements  in 
the  solid  or  liquid  state. 

The  Quantum  Theory.  Quite  recently  a  new  theory  of 
molecular,  atomic  and  radiant  energy,  called  the  quantum  theory, 
has  been  developed  by  Plank,  Einstein,  Nernst,  Sackur,  Debye, 
Sommerfeld  and  others.  The  theory  supposes  that  there  are  in 
the  atoms  of  the  elements,  or  associated  with  them,  resonators  or 
oscillators  of  such  a  nature  that  they  can  emit  energy  only  in 
definite,  unit  quantities.  The  resonators  may  be  atoms,  ions,  or 
electrons ;  i.e.  they  may  be  particles  with  or  without  electrical 
charges.  The  theory  seems  to  give  a  satisfactory  explanation  of 
the  low  and  variable  specific  heats  of  some  of  the  elements,  of 
some  photo-electric  effects  which  were  previously  hard  to  under- 
stand and  of  a  variety  of  other  phenomena.  From  the  character 
of  the  men  who  are  working  on  the  theory  and  the  striking 
results  already  attained  it  seems  likely  to  be  developed  very 
rapidly  in  the  near  future. 

Sodium,  Na,  23.  Sodium  chloride,  NaCl,  or  common  salt, 
forms  about  75  per  cent  of  the  residue  left  by  the  evaporation 
of  sea  water.  It  is  also  found  in  enormous  beds  of  rock  salt  in 
Germany,  Louisiana,  Kansas,  Utah  and  elsewhere  and  in 
strong  brines  found  by  boring  deep  wells  in  very  many  places. 
Sodium  is  taken  up  by  seaweeds  in  their  growth  very  much  as 
potassium  is  taken  up  by  land  plants,  and  the  ashes  of  sea- 
weeds contain  considerable  quantities  of  sodium  carbonate. 
Deposits  of  sodium  sesquicarbonate,  NaHCO3.Na2CO3.2H2O, 
called  trona,  of  sufficient  extent  to  be  of  industrial  importance 
are  found  in  Egypt  and  in  Venezuela.  The  occurrence  of 


SODIUM  399 

borax,  Na2B4O7.10H2O,  in  lakes  in  California  has  been  men- 
tioned; also  that  of  sodium  nitrate,  NaNOs.  These  are  of 
value,  primarily,  for  the  boron  and  nitrogen  which  they  contain. 
Sodium  is  a  constituent  of  practically  all  silicious  rocks. 

Metallurgy.  Properties.  Metallic  sodium  and  potassium 
were  first  prepared  by  Sir  Humphry  Davy  in  London  in  1807 
by  the  electrolysis  of  moist  sodium  and  potassium  hydroxides. 
The  discovery  awakened  very  great  interest,  both  because  the 
metals  showed  very  striking  properties,  quite  different  from 
those  of  any  metals  hitherto  known,  and  because  it  indicated 
very  clearly  that  many  other  earthy  substances  contain  ele- 
ments which  could  not  at  that  time  be  prepared  in  the  free 
state.  Within  a  few  years  the  new  metals  proved  effective 
agents  for  the  liberation  of  a  number  of  other  elements. 

Sodium  may  also  be  prepared  by  heating  a  mixture  of  sodium 
carbonate  and  carbon : 

Na2CO3  +  2C  =  3CO  +  2Na 

The  sodium,  which  boils  at  742°,  distills  from  the  mixture  and 
is  collected  in  iron  condensers.  During  comparatively  recent 
years  metallic  sodium  is  prepared  commercially  by  various 
electrolytic  methods,  from  the  hydroxide,  the  nitrate  or  the 
chloride. 

Sodium  is  a  silver  white  metal  which  tarnishes  instantly  on 
exposure  to  moist  air.  In  dry  air  at  300°  to  400°  it  is  oxidized 
to  sodium  peroxide,  Na2O2.  It  is  kept  in  sealed  cans  or  under 
kerosene  to  protect  it  from  the  action  of  the  air.  It  melts  at 
97.5°  and  boils  at  742°,  giving  a  dark  green  vapor.  The  specific 
gravity  of  the  solid  is  0.97.  When  a  small  piece  of  sodium  is 
thrown  on  water  the  heat  of  the  reaction  causes  it  to  melt  and 
the  globule  of  metal  rolls  rapidly  over  the  surface  of  the  water 
without  taking  fire  or  igniting  the  hydrogen,  differing  in  this 
respect  from  potassium.  If  the  metal  is  thrown  on  a  piece  of 
filter  paper  floating  on  the  water,  the  heat  is  concentrated  and 
the  hydrogen  takes  fire  and  burns  with  a  yellow  flame.  In 
both  cases,  of  course,  sodium  hydroxide  is  formed. 


400  A  TEXTBOOK  OF  CHEMISTRY 

Sodium  is  manufactured  in  considerable  amounts  for  use  in 
preparing  sodium  peroxide,  Na2O2,  for  the  preparation  of  a 
mixture  of  potassium  and  sodium  cyanides  from  potassium  ferro- 
cyanide  (p.  319),  and  for  use  in  the  synthesis  of  a  variety  of 
organic  compounds. 

The  Alkali  Industry.  Sodium  is  an  essential  constituent  of 
common  soap,  of  glass,  of  salsoda,  or  washing  soda,  and  of 
baking  soda.  As  common  salt,  NaCl,  is  much  cheaper  than 
any  other  compound  of  sodium,  it  now  furnishes  the  basis  for 
the  preparation  of  all  of  these  substances,  but  it  is  necessary  to 
prepare  from  it,  at  first,  one  of  the  sodium  carbonates  or  sodium 
hydroxide.  Till  the  close  of  the  eighteenth  century  an  impure 
sodium  carbonate  from  Egypt  and  the  ash  of  seaweeds  were 
used  as  the  sources  of  sodium  carbonate  and  sodium  hydroxide 
for  the  manufacture  of  hard  soap,  while  potassium  carbonate 
from  wood  ashes  was  also  extensively  used  for  the  manufacture 
of  soft  soap.  During  the  disturbed  commercial  relations  which 
followed  the  French  Revolution,  the  foreign  supplies  of  sodium 
carbonate  were  cut  off  and  all  available  potassium  compounds 
were  needed  for  the  manufacture  of  gunpowder.  This  caused* 
the  French  government  to  offer  a  prize  for  a  method  of  manu- 
facturing sodium  carbonate  from  salt.  The  prize  was  awarded 
to  Leblanc  and  his  process  was  used  for  a  short  time  in  France, 
but  could  not  there  compete  with  the  sodium  carbonate  from 
other  sources  when  commercial  relations  with  other  countries 
were  again  established.  Leblanc  himself  did  not  secure  any 
permanent  advantage  from  his  invention  and  died  in  a  poor- 
house.  About  twenty  years  later  Musgrave,  in  England,  took 
up  the  process  again  and  succeeded  in  making  it  a  commercial 
success.  It  held  the  field  of  alkali  manufacture  almost  without 
competition  for  fifty  years.  About  1860  it  had  to  meet  the 
competition  of  the  ammonia-soda  process,  the  principles  of 
which  had  been  discovered  in  1838,  but  which  was  first  put 
into  successful  operation  by  Solvay  more  than  twenty  years 
later.  From  then  till  the  close  of  the  nineteenth  century  the 
Leblanc  process  continually  lost  ground  in  competition  with  tke 


SODIUM  HYDROXIDE  401 

Solvay  manufacture,  maintaining  a  precarious  existence  only  by 
the  most  careful  conservation  of  the  by-products,  hydrochloric 
acid  or  chlorine  and  sulfur.  In  1900  only  two  large  Leblanc  fac- 
tories remained  in  the  world,  one  in  England  and  one  in  Germany. 

The  most  serious  difficulty  with  the  ammonia-soda  process  is 
that  the  chlorine  of  the  salt  is  converted  into  calcium  chloride 
or  magnesium  chloride,  practically  worthless  products.  In  the 
closing  years  of  the  nineteenth  century,  with  the  stimulus  of 
comparatively  cheap  electrical  energy,  many  electrolytic  pro- 
cesses were  developed  which  give  both  chlorine  and  sodium 
hydroxide  directly  from  salt.  It  seems  probable  that  these 
processes  will  eventually  displace  the  Solvay  process,  at  least 
for  the  production  of  caustic  alkali. 

Sodium  Hydroxide.  So  long  as  trona  from  Egypt  or  the  ashes 
of  sea  plants  were  used,  sodium  hydroxide  was  prepared  by 
treating  a  solution  of  these  with  'slaked  lime.  A  more  pure 
sodium  hydroxide  was  prepared  in  the  same  way  from  the 
sodium  carbonate  of  the  Leblanc  or  ammonia-soda  processes : 

Na2CO3  +  Ca(OH)2  =  2  NaOH  +  CaCO3 

The  reaction  depends,  of  course,  on  the  insolubility  of  the 
calcium  carbonate.  In  very  concentrated  solutions  the  reaction 
may  be  reversed  because  calcium  hydroxide,  Ca(OH)2,  is  also 
difficultly  soluble,  and  with  a  high  concentration  of  hydroxide 
ions,  OH~,  the  solubility  product  for  that  substance  may  be 
exceeded  even  in  a  solution  containing  so  insoluble  a  salt  as 
calcium  carbonate. 

As  has  been  stated  above,  sodium  hydroxide  is  now  prepared 
on  a  large  scale  by  electrolysis.  Many  different  patents  have 
been  issued  for  such  processes  and  it  is  probably  too  soon  to 
decide  which  forms  are  likely  to  prove  permanently  suitable. 
In  some  forms  a  diaphragm,  usually  of  asbestos,  is  used  to  sepa- 
rate the  anode  space  from  the  cathode.  The  anode  must  be  of 
carbon  or  of  platinum  or  platinum-iridium  and  the  anode  space 
is  inclosed  so  that  the  chlorine  liberated  may  be  collected  and 
utilized.  The  cathode  is  usually  of  iron.  Chlorine  is  liberated 


402 


A  TEXTBOOK  OF  CHEMISTRY 


at  the  anode,  while  hydrogen  is  liberated  at  the  cathode  and 
sodium,  Na+,  and  hydroxide,  OH~,  ions  remain  in  solution, 
the  hydrogen,  of  course,  coming  from  the  water,  though  the 
transfer  of  ions  through  the  solution  is  mainly  that  of  sodium 
and  chloride  ions.  The  points  aimed  at  are  to  secure  as  high  a 
concentration  of  hydroxide,  OH~,  and  as  low  a  concentration 
of  chloride,  Cl~,  ions  as  possible  in  the  cathode  space  and  the 
reverse  of  this  around  the  anode.  To  this  end  the  salt  solution 
is  continuously  introduced  at  the  anode  while  the  hydroxide 
solution  is  removed  from  the  cathode.  The  presence  of  hy- 
droxide at  the  anode  leads  to  the  formation  of  hypochlorite  and 
loss  of  current.  The  hydroxide  solution  may  be  concentrated 
till  nearly  all  of  the  salt,  NaCl,  remaining  in  it  separates,  leav- 
ing a  solution  in  which  nearly  all  of  the  sodium  is  in  the  form  of 
hydroxide.  This  solution  is  then  evaporated  till  the  water  has 
been  expelled,  which  requires  a  comparatively  high  temperature. 
The  sodium  hydroxide  obtained  in  this  way  is  sufficiently  pure 
for  the  manufacture  of  soap  and  for  many  other  industrial  uses. 
The  Castner-Kellner  apparatus  gives  an  almost  pure  solution 
of  sodium  hydroxide  directly.  It  consists  of  a  slate  box  divided 

into  three  com- 
partments by  two 
partitions,  which 
fit  only  loosely  in 
grooves  in  the 
bottom  of  the  box 
(Fig.  94).  Mer- 


.JL 


Fig.  94 


cury  placed  on  the  bottom  of  the  box  seals  these,  giving  a  contin- 
uous metallic  layer  for  the  three  compartments,  but  prevents  a 
dilute  solution  of  sodium  hydroxide  placed  in  the  central  com- 
partment from  mixing  with  the  brine  placed  in  the  two  side 
compartments.  Graphite  anodes  are  placed  in  the  two  side 
compartments  and  an  iron  cathode  in  the  central  one.  Chlo- 
rine is  evolved  from  the  anodes  and  is,  of  course,  collected  and 
used  for  the  manufacture  of '  bleaching  powder  or  for  some 
other  purpose.  The  mercury  in  the  two  side  compartments  is 


SODIUM  HYDROXIDE 


403 


negative  as  compared  with  the  graphite  anodes  and  the  sodium 
liberated  at  its  surface  combines  with  it  to  form  a  liquid  sodium 
amalgam.  By  a  slight  tilting  motion  the  amalgam  is  caused  to 
flow  alternately  to  one  side  or  the  other  and  so  is  brought 
into  the  central  compartment.  Here  it  is  positive  with  refer- 
ence to  the  more  negative  cathode  and  the  hydroxide  ions 
brought  to  its  surface  by  the  current  combine  with  the  sodium 
of  the  amalgam  to  form  sodium  hydroxide,  while  the  hydrogen 
ions  of  the  water  are  discharged  and  liberated  as  free  hydrogen, 
H2,  at  the  surface  of  the  iron  cathode.  The  hydroxide  solu- 
tion is  kept  at  a  constant  concentration  by  introducing  water 
at  one  side  and  removing  some  of  the  solution  at  the  other. 
Salt  is  added  from  time  to  time  to  the  side  compartments. 

Sodium  hydroxide  is  a  white  solid,  which  melts  at  a  red  heat* 
It  deliquesces  on  exposure  to  the  air,  but  the  solution  soon  ab- 
sorbs carbon  dioxide  and  then  evaporates,  leaving  a  residue  of 
sodium  carbonate. 

Sodium  hydroxide  is  used  in  the  manufacture  of  ordinary 
hard  soaps  and  may  be  used  for  the  preparation  of  many  of 
the  salts  of  sodium. 

The  specific  gravity  of  solutions  of  different  concentrations 
is  as  follows : 

DENSITY  OF  SOLUTIONS  OF  SODIUM  HYDROXIDE 


SPECIFIC  GRAVITY 

PER  CENT 
NaOH 

GRAMS  OP  NaOH 

IN   100  CO. 

1.0555 

5 

5.277 

1.1111 

10 

11.111 

1.1665 

15 

17.497 

1.2219 

20 

24.438 

1.2771 

25 

31.928 

1.3312 

30 

39.936 

1.3838 

35 

48.433 

1.4343 

40 

57.372 

1.4828 

45 

66.726 

1.5303 

50 

76.515 

404  A  TEXTBOOK  OF  CHEMISTRY 

Sodium  hydroxide  dissolves  in  water  with  the  evolution  of 
considerable  heat,  and  it  will  be  seen  from  the  table  that  the 
addition  of  a  small  amount  of  sodium  hydroxide  causes  the 
volume  of  the  water  to  diminish. 

Sodium  Oxide,  Na2O.  With  the  exception  of  lithium  oxide, 
Li2O,  the  oxides  of  the  alkali  metals  cannot  be  prepared  by 
heating  the  hydroxides  or  carbonates.  In  this  respect  they 
differ  from  all  other  metallic  oxides.  Sodium  oxide  may  be 
prepared  by  heating  sodium  hydroxide,  NaOH,  with  metallic 
sodium.  It  combines  with  water  to  form  the  hydroxide,  but 
has,  at  present,  no  practical  importance. 

Sodium  Peroxide,  Na2O2,  is  prepared  by  heating  metallic 
sodium  to  300°^400°  in  dry  air.  The  sodium  is  placed  in  shallow 
aluminium  trays,  which  are  passed  slowly  through  long  ovens 
one  way,  while  air  passes  in  the  opposite  direction.  In  this 
way  the  pure  sodium  comes  at  first  in  contact  with  air  which 
has  been  deprived  of  most  of  its  oxygen  and  a  too  vigorous 
action  is  avoided,  while  the  action  is  finally  completed  at  the 
other  end  with  fresh  air. 

The  use  of  fused  sodium  peroxide  containing  a  little  copper 
oxide,  under  the  name  of  "  oxone,"  for  the  preparation  of 
oxygen  has  been  mentioned  (p.  21).  It  may  be  converted  by 
cold,  moist  air  into  a  hydrate,  Na2O2.H2O,  which  can  be  dis- 
solved in  water  with  little  decomposition.  It  is  hydrolyzed,  how- 
ever, to  sodium  hydroxide,  NaOH,  and  hydrogen  peroxide,  H2O2. 

On  treatment  with  cold,  dilute  acids  sodium  peroxide  gives  a 
solution  of  hydrogen  peroxide,  H2O2,  which  is  used  to  bleach 
silk,  wool,  hair  and  other  substances  which  would  be  affected 
injuriously  by  chlorine. 

Sodium  peroxide  is  also  a  very  valuable  oxidizing  agent  for 
many  laboratory  uses. 

Sodium  Chloride.  Salt  is  sometimes  obtained  by  direct 
mining,  but  rock  salt  is  rarely  sufficiently  pure  for  direct  use,  and 
it  offers  especial  difficulties  in  mining,  owing  to  its  effect  in 
dulling  tools  used  to  cut  it  and*  because  blasting  does  not  loosen 
it  up  satisfactorily.  It  is  found  better  to  prepare  an  opening 


SODIUM  CHLORIDE 


405 


in  the  bed  of  salt  and  allow  water  to  stand  in  contact  with  it 
till  a  saturated  solution  is  obtained,  many  of  the  impurities 
present  remaining  undissolved  and  settling  to  the  bottom. 
The  solution  is  then  pumped  out  and  evaporated  to  crystallize 
the  salt.  As  salt  is  nearly  as  soluble  in  cold  as  in  hot  water,  it 
cannot  be  crystallized  practically  by  cooling  a  hot  solution. 
In  some  places,  especially  at  Syracuse,  New  York,  and  in  Michi- 
gan, saturated  brines  are  obtained  directly  from  artesian  wells. 
For  the  evaporation  of  the  brines  triple-effect  evaporators 
are  used  to  advantage.  The  principle  of  these  is  shown  in  the 
accompanying  diagram  (Fig.  95).  The  brine  in  pan  A  is  heated 
directly,  or  by  superheated  steam  in  coils  or  in  a  false  bottom, 


Fig.  95 

and  it  boils  under  atmospheric  pressure.  The  steam  from  this 
pan  passes  under  B,  in  which  a  pressure  of  perhaps  550  mm.  is 
maintained  so  that  the  condensation  of  the  steam  from  the  first 
pan  beneath  it  will  cause  the  brine  which  it  contains  to  boil. 
The  steam  from  this  will,  in  turn,  cause  the  brine  in  C  to  boil 
under  a  pressure  of  300  mm.  By  such  an  arrangement  a  given 
weight  of  coal  will  evaporate  nearly  three  times  as  much  water 
as  it  would  if  used  directly  in  the  usual  manner.1 

Commercial  salt  contains  small  quantities  of  various  impuri- 
ties. The  most  objectionable,  perhaps,  is  magnesium  chloride, 

1  It  may  be  remarked,  incidentally,  that  when  this  process  is 
used  for  the  preparation  of  distilled  water,  from  waters  containing 
relatively  small  amounts  of  solids  in  solution  and  the  successive 
differences  of  pressure  may  be  much  less,  as  many  as  ten  boilers 
may  be  used  in  series.  For  a  description  of  the  Yaryan  Evapora- 
tor, which  uses  a  modification  of  this  system,  see  J.  Soc.  Chem. 
Ind.  14,  112  (1895). 


406  A  TEXTBOOK  OF   CHEMISTRY 

which  makes  it  hygroscopic  or  even  deliquescent  in  moist  air. 
Pure  sodium  chloride  can  be  obtained  by  precipitating  a  solu- 
tion of  salt  with  concentrated  hydrochloric  acid. 

Sodium  chloride  crystallizes  in  cubes.  It  melts  at  820°  and 
may  be  volatilized  at  a  high  temperature.  The  crystals  usually 
decrepitate  on  heating,  owing  to  water  inclosed  in  them. 

Salt  is  an  essential  constituent  of  human  diet,  furnishing 
chlorine  for  the  hydrochloric  acid  of  the  gastric  juice. 

Sodium  Sulfate.  Glauber's  Salt,  Na2SO4.10  H2O.  By  heat- 
ing salt  with  the  theoretical  amount  of  sulfuric  acid  it  may  be 
converted  almost  quantitatively  into  anhydrous  sodium  sulfate, 
Na2SO4.  The  operation  is  carried  out  on  a  large  scale  as  the 
first  step  in  the  Leblanc  soda  process.  The  anhydrous  sulfate 
is  also  used  in  the  manufacture  of  glass. 

The  crystallized  hydrate,  Na2SO4.10  H2O,  has  a  solubility 
in  water  which  increases  very  rapidly  with  rising  temperature 
till  the  transition  point,  32.383°,  is  reached.  If  crystals  of  the 
hydrate  are  heated  above  this  temperature,  they  are  transformed 
into  a  mixture  of  anhydrous  sodium  sulfate  and  a  saturated 
solution  of  the  latter,  which  is  less  soluble  than  the  crystallized 
hydrate  at  temperatures  above  this  point.  The  transition  is 
accompanied  by  an  absorption  of  heat  in  very  much  the  same 
manner  as  the  melting  of  ice  and  may  be  used  as  an  accurate, 
fixed  point  for  the  correction  of  thermometers.  (Richards,  Z. 
physik.  Chem.  43,  465.) 

*  The  transition  point  is  a  quadruple  point  in  the  nomenclature 
of  the  phase  rule  (p.  107),  the  four  phases  being  water  vapor  at 
a  pressure  of  30.8  mm.,  the  hydrate,  Na2SO4.10H2O,  the 
anhydrous  salt,  Na2SO4,  and  the  saturated  solution.  As  there 
are  four  phases  and  only  two  components,  sodium  sulfate  and 
water,  the  system  is  invariant  and  there  can  be  no  change  in 
temperature  or  pressure  without  the  disappearance  of  one  of  the 
phases.1 

1  Practically,  the  transition  point  is  determined  in  contact  with 
air  at  atmospheric  pressure  and  the  effect  of  pressure  is  not  con- 
sidered. In  the  presence  of  the  vapor  phase  only,  the  temperature 


SODIUM  SULFATE 


407 


The  solubility  of  sodium  sulfate  is  shown  in  the  diagram,  Fig. 
96.  The  concentrations  are  given  for  the  anhydrous  salt  through- 
out. Below  the  transition  point  the  hydrate  separates  on  evap- 
orating or  cooling  the  solution,  though  supersaturated  solutions 


Solubility— Grams  of  Na2SO4in  100  grams  of  water 

1-1  o  o  o  o  o  o 


c 

' 

S 

oo 
ta 

\ 

\ 

o 

0 

eg 

X. 

/ 

0 
c 

a 

--*J 

§0 

o 

§ 

M 
C 

/ 

°°  Fig.  96 

of  the  transition  point  would  be  slightly  different,  just  as  the  true 
transition  point  for  water,  ice,  vapor  is  0.0076°  higher  than  the  melt- 
ing point  of  ice  under  atmospheric  pressure,  which  is  used  as  the 
zero  point  for  thermometers. 


408  A  TEXTBOOK  OF  CHEMISTRY 

are  easily  formed  (p.  80).  Above  the  transition  point  the  anhy- 
drous salt  separates  on  heating  or  on  evaporation. 

Sodium  sulfate  is  one  of  the  chief  active  constituents  of 
Hunyadi  water  and  of  some  other  similar  medicinal  waters. 

Acid  Sodium  Sulfate  or  Sodium  Bisulfate,  NaHSO4,  is  formed 
as  the  first  step  in  the  preparation  of  the  sulfate  from  salt,  or 
in  the  manufacture  of  nitric  acid  from  sodium  nitrate,  NaNO3o 
At  about  300°  it  loses  water  and  is  converted  into  sodium  pyro- 
sulfate,  Na2S2O7.  At  a  still  higher  temperature  it  is  decom- 
posed into  sodium  sulfate  and  sulfur  trioxide,  SO3.  Sodium 
pyrosulfate  is  used  in  analytical  chemistry  to  dissolve  aluminium 
oxide,  A12O3,  ferric  oxide,  Fe2O3,  and  titanium  oxide,  TiO2,  after 
these  have  been  brought  to  a  difficultly  soluble  form  by  ignition. 
(Hillebrand,  Analysis  of  Silicate  and  Carbonate  Rocks,  Bulletin 
422,  U.  S.  Geol.  Survey,  p.  105.) 

Sodium  Sulfite,  Na2SO3.H2O,  is  prepared  by  burning  sulfur 
and  passing  the  sulfur  dioxide  formed  through  a  solution  of 
sodium  carbonate.  It  is  used  as  a  reducing  agent  in  photog- 
raphy and  for  the  preparation  of  sodium  thiosulfate. 

Acid  Sodium  Sulfite,  or  Sodium  Bisulfite,  NaHSO3,  is  pre- 
pared by  passing  sulfur  dioxide  in  excess  into  a  solution  of  sodium 
carbonate.  It  is  sometimes  used  as  an  addition  to  cider  to  stop 
fermentation.  A  40  per  cent  solution  is  very  conveniently 
used  for  the  preparation  of  sulfur  dioxide  in  the  laboratory. 

Sodium  Hyposulfite,  Na2S2O4.  A  solution  of  this  salt  is 
prepared  by  the  action  of  zinc  on  a  solution  of  sodium  bisulfite, 
NaHSO3,  containing  an  excess  of  sulfurous  acidjHaSOs. 

2  NaHSO3  +  H2S03  +  Zn  =  Na2S2O4  +  ZnSO3  +  2  H2O 

Sodium  hyposulfite  is  a  powerful  reducing  agent  and  is  used 
especially  to  reduce  indigo  to  indigo  white  (p.  341). 

Sodium  Thiosulfate,  Na2S2O3.5  H2O,  is  prepared  by  dissolv- 
ing sulfur  in  a  solution  of  sodium  sulfite.  It  dissolves  the  halides 
of  silver  (AgCl,  AgBr  and  Agl)  and  is  used  to  fix  photographic 
pictures  (p.  445).  It  is  usually  called  by  photographers  and 
pharmacists  "  sodium  hyposulfite."  It  has  also  been  used  in 


SODIUM  SULFIDE  409 

extraction  of  silver  from  its  ores  in  the  so-called  "  hyposulfite- 
lixiviation  "  processes. 

The  anhydrous  thiosulfate,  Na2S2O3,  can  be  obtained  by  dry- 
ing the  crystals  at  a  moderate  temperature.  If  the  dry  salt  is 
heated,  it  decomposes  into  sodium  sulfate,  sodium  sulfide,  and 
sulfur  or  sodium  polysulfide. 

If  a  solution  of  sodium  thiosulfate  is  warmed  with  copper 
sulfate,  cuprous  sulfide,  Cu2S,  and  sulfur  are  precipitated,  while 
sodium  sulfate  remains  in  solution.  These  reactions  show  that 
the  compound  retains  the  reducing  properties  of  the  sulfites 
and  also  the  properties,  in  part,  of  a  sulfide.  This  recalls  the 
method  of  preparation  and  agrees  well  with  the  formula 
Na— Sv  ^O 

/Sx'    ,  which  is  assigned  to  the  compound. 
Na— CK     ^O 

Sodium  Tetrathionate,  Na2S4O6.H2O,  is  formed  by  the  action 
of  iodine  on  a  solution  of  sodium  thiosulfate : 

2  Na2S2O3  +  I2  =  Na2S2O6  +  2  Nal 

The  reaction  is  quantitative  and  is  much  used  in  volumetric 
analysis. 

Sodium  Sulfide,  Na2S,  may  be  prepared  by  passing  hydrogen 
sulfide  into  a  solution  of  sodium  hydroxide  in  the  requisite 
amount  and  evaporating  the  solution  to  dryness  with  exclusion 
of  air.  It  is  hydrolyzed  by  water,  giving  a  strongly  alkaline 
solution. 

If  a  mixture  of  sodium  carbonate  and  sulfur  is  heated,  or  if 
any  metallic  sulfide  is  heated  with  sodium  carbonate  on  charcoal 
or  if  any  metallic  sulfate  is  heated  with  sodium  carbonate  in  the 
reducing  flame  (p.  304)  on  charcoal,  a  sulfide  is  formed.  If  this 
is  moistened  with  water  on  a  silver  coin  a  black  spot  of  silver 
sulfide,  Ag2S,  will  be  formed.  The  reaction  is  used  as  a  test 
for  sulfur  in  any  form  of  inorganic  combination. 

Sodium  Hydrosulfide,  NaSH,  is  formed  when  hydrogen  sulfide 
is  passed  into  a  solution  of  sodium  hydroxide  in  twice  the  amount 
necessary  to  form  the  sulfide.  It  loses  hydrogen  sulfide  on  evap- 


410  A  TEXTBOOK  OF  CHEMISTRY 

oration  of  the  solution  with  exclusion  of  air,  in  the  same  way  that 
sodium  bicarbonate,  NaHCOs,  loses  carbon  dioxide. 

Sodium  Nitrate,  NaNO3,  is  found  in  immense  beds  in  Chile, 
South  America.  It  has  been  the  chief  source  from  which  nitric 
acid  and  saltpeter,  KNO3,  have  been  prepared  and  large  quan- 
tities have  also  been  used  for  fertilizers,  to  furnish  the  nitrogen 
necessary  for  the  growth  of  crops.  It  is  estimated  that  the 
supply  from  Chile  will  be  exhausted  in  a  comparatively  few 
years,  but  there  seems  now  a  good  probability  that  the  manu- 
facture of  nitrates  from  the  nitrogen  of  the  air  will  soon  be  in  a 
position  to  supply  its  place. 

Sodium  nitrate  crystallizes  without  water  of  crystallization 
in  rhombohedra.  It  melts  at  316°.  It  is  hygroscopic  and  for 
that  reason  cannot  be  used  in  place  of  potassium  nitrate  for  the 
manufacture  of  ordinary  gunpowder,  though  its  low  molecular 
weight  makes  it,  otherwise,  more  suitable. 

Sodium  Nitrite,  NaNO2,  is  prepared  by  heating  sodium  nitrate 
with  metallic  lead.  It  is  very  easily  soluble,  but  crystallizes 
well.  It  is  used  in  laboratories  and  in  factories  for  the  prepara- 
tion of  diazonium  compounds  for  the  manufacture  of  dyestuffs 
and  other  important  compounds. 

Sodamide,  NaNH2,  is  prepared  by  passing  ammonia  over  me- 
tallic sodium  at  300°-350°.  (See  Dennis  and  Browne,  J.  Am. 
Ch.  Soc.  26,  587.)  It  dissolves  in  liquid  ammonia,  ionizing  to 
Na+  and  NH2~  (p.  207).  It  is  hydrolyzed  by  water  to  sodium 
hydroxide  and  ammonia.  It  has  recently  become  important 
for  the  preparation  of  indigo. 

*  Sodium  Trinitride,  NaN3,  is  formed  by  the  action  of  nitrous 
oxide,  N2O,  on  sodamide  (p.  223). 

Disodium  Phosphate,  Na2HPO4.12  H2O,  is  the  best  known 
and  most  important  of  the  phosphates  of  sodium.  It  is  isomor- 
phous  with  the  corresponding  arsenate,  Na2HAsQ4-12  H2O, 
and  usually  contains  some  of  that  salt  derived  from  the  impure 
phosphorus  or  from  the  sulfuric  acid  used  in  the  preparation  of 
phosphoric  acid.  The  salt  is  sometimes  used  in  medicine  as 
a  mild  cathartic. 


SODIUM  CARBONATE  411 

Sodium  Carbonate  or  Sal  soda  (Washing  Soda), 
Na2CO3.10  H2O.  The  Leblanc  Soda  Process.  The  Leblanc 
soda  process  is  carried  out  in  three  operations. 

1.  Salt  is  treated  with  sulfuric  acid  on  the  hearth  of  a  furnace 
which  can  be  heated  to  complete  the  reaction  and  expel  all  of 
the  hydrochloric  acid.  The  latter  is  conveyed  through  a  tower 
filled  with  coke  over  which  water  is  trickling  and  the  aqueous 
hydrochloric  acid  obtained  is  sold  or  used  for  the  preparation  of 
chlorine.  In  the  early  years  of  manufacture  the  acid  was  al- 
lowed to  escape  and  produced  disastrous  effects  upon  vegetation 
in  the  neighborhood  of  the  works.  This  led  to  stringent  legis- 
lation forbidding  the  escape  of  the  acid.  Later,  the  recovery 
of  the  hydrochloric  acid  proved  profitable,  and  this  has  been 
an  important  factor  in  preventing  the  complete  abandonment 
of  the  process : 

12  NaCl  +  H2S04  =  Na2SO4  +  2  HC1 
2.  The  sodium  sulfate  is  mixed  with  charcoal,  or  coal,  and 
limestone,  CaCO3,  and  heated  to  fusion  in  the  "  black  ash  fur- 
nace." The  mass  melts,  the  sodium  sulfate  is  reduced  to  sodium 
sulfide,  Na2S,  and  the  latter  reacts  with  the  calcium  carbonate 
to  form  calcium  sulfide,  CaS,  and  sodium  carbonate : 

Na2S04  +  2  C  =  Na2S  +  2  CO2 
Na2S  +  CaCO3  =  CaS  +  Na2CO3 

3.  The  "  black  ash,"  when  cold,  is  leached  with  water,  which 
dissolves  the  sodium  carbonate  and  leaves  most  of  the  calcium 
sulfide  undissolved.  There  is  a  tendency  for  the  calcium  sulfide 
to  hydrolyze  to  calcium  hydrosulfide,  Ca(SH)2,  and  calcium 
hydroxide,  Ca(OH)2,  but  this  takes  place  slowly  and  is  repressed 
in  the  strongly  alkaline  solution,  so  that  a  fairly  complete  separa- 
tion is  obtained.  The  sodium  carbonate  is  obtained  from  the 
solution  by  evaporation  and  crystallization.  If  the  evaporation 
and  crystallization  take  place  above  35.2°,  the  monohydrate, 
Na2CO3.H2O,  crystallizes  from  the  solution.  If  the  salt  is  crys- 
tallized below  that  temperature,  by  evaporation  or  by  cooling 


412  A  TEXTBOOK  OF  CHEMISTRY 

the  solution,  the  dekahydrate,  Na2CO3.10  H2O,  is  formed. 
The  temperature  given,  35.2°,  is,  of  course,  the  transition  point 
from  the  dekahydrate  to  the  monohydrate. 

The  Leblanc  process,  which  was  so  important  during  the 
nineteenth  century,  seems  likely  to  be  completely  displaced  by 
other  processes  in  the  near  future. 

Crystallized  sodium  carbonate  is  known  as  sal  soda,  or  washing 
soda,  and  is  used  for  laundry  purposes,  for  the  manufacture  of 
soap  and  for  softening  water.  Anhydrous,  or  calcined,  sodium 
carbonate  is  used  in  the  manufacture  of  glass  and  for  the  pre- 
paration of  other  compounds  of  sodium. 

Sodium  carbonate  is  hydrolyzed  by  water  to  sodium  bicar- 
bonate and  sodium  hydroxide,  owing  to  the  very  trifling  ioniza- 
tion  of  the  hydrocarbonate  ion,  HCO3~  : 

Na+  +  Na+  +  CO3=  +  H+  +  OH~ 

=  Na+  +  Na+  +  HCO3~  +  OH~ 

Sodium  Bicarbonate,  or  Baking  Soda,  NaHCO3.  The  Ammo- 
nia-soda Process.  About  1860  Solvay  succeeded  in  putting  this 
process,  which  had  been  discovered  many  years  before,  into  a 
successful  form.  It  has  been  found  more  economical  than  the 
Leblanc  process  and  by  the  close  of  the  nineteenth  century  it  had 
very  largely  displaced  that  method  of  manufacture. 

A  strong  brine  is  first  treated  with  ammonia  till  it  contains 
one  molecule  for  each  molecule  of  salt.  A  tower  about  20  meters 
high  and  having  a  series  of  shelves  is  filled  with  the  brine  and 
carbon  dioxide  is  forced  into  the  tower  at  the  bottom.  Am- 
monium bicarbonate,  NH4HCO3,  is  at  first  formed,  but  as  sodium 
bicarbonate,  NaHCO3,  is  the  least  soluble  of  the  various  com- 
binations of  ions  possible,  it  separates  and  is  deposited  in  layers 
on  the  shelves : 

NH3  +  H2O  +  CO2  =  NH4HCO3 
NH4HC03  +  NaCl  =  NH4C1  +  NaHCO3 

The  sodium  bicarbonate,  after  removal  from  the  shelves  and 
washing  in  centrifugals  with  a  little  water,  is  almost  chemically 


SODIUM  SILICATE  413 

pure.     Sodium  carbonate  can  be  easily  obtained  from  this  by 
calcining  it  at  a  comparatively  low  temperature : 

2  NaHCO3  =  Na2CO3  +  CO2  +  H2O 

A  large  part  of  the  bicarbonate  is  converted  to  carbonate  for 
the  market  by  this  process.  The  carbon  dioxide  obtained  is,  of 
course,  returned  to  use  in  the  first  stage  of  the  manufacture. 

Ammonia  is  much  more  valuable  than  sodium  carbonate,  at 
the  present  time,  and  the  ammonia-soda  process  depends, 
economically,  upon  the  recovery  of  the  ammonia.  This  is  ef- 
fected by  treating  the  solution  of  ammonium  chloride  with 
slaked  lime,  Ca(OH)2,  and  distilling : 

2  NH4C1  +  Ca(OH)2  =  CaCl2  -f  2  NH3 

In  well-conducted  factories  not  more  than  5  kilos  of  ammonia 
are  lost  in  preparing  1000  kilos  of  sodium  carbonate. 

Sometimes  magnesium  oxide,  MgO,  is  used  in  place  of  calcium 

hydroxide,  as  it  is  possible  to  recover  chlorine  or  hydrochloric 

'  acid  from  the  magnesium  chloride,  but  the  fact  that  the  chlorine 

is  left  by  the  process  in  the  form  of  a  comparatively  worthless 

compound  may  lead,  ultimately,  to  its  abandonment. 

Sodium  bicarbonate  is  used  as  "  baking  soda  "  in  cooking,  to 
furnish  carbon  dioxide  for  lightening  bread  or  cake.  It  is  used 
with  sour  milk  or,  frequently,  mixed  with  cream  of  tartar,  tar- 
taric  acid,  alum  or  acid  calcium  phosphate  in  the  various  baking 
powders.  It  is  also  used  as  a  mild  alkali  in  medicine. 

Sodium  Silicate  or  Soluble  Glass,  Na2SiO3,  is  made  by  fusing 
sodium  carbonate  and  sand  in  the  proper  proportion.  It  is 
hydrolyzed  by  water  and  the  solution  reacts  strongly  alkaline, 
but  the  silicic  acid  remains  in  colloidal  solution.  The  solution 
is  used  to  fireproof  wood  and  fabrics,  covering  them  with  a  thin, 
glassy  coating,  which  renders  them  much  less  inflammable. 

Sodium  Tetraborate  or  Borax,  Na2B4O7.10  H2O,  has  been  con- 
sidered in  connection  with  boron  (p.  367). 


CHAPTER  XXIV 

ALKALI     METALS     (Continued)  :     POTASSIUM,     AMMONIUM, 
RUBIDIUM,    CAESIUM :  THE    SPECTROSCOPE 

Potassium,  K,  39.10.  Occurrence.  Many  of  the  natural  sili- 
cates and  especially  the  potash  feldspar,  orthoclase,  KAlSisOg, 
contain  potassium.  These  feldspars  form  an  essential  constitu- 
ent of  granite  and  in  the  disintegration  of  granites  and  other 
rocks  through  the  action  of  water  during  very  long  periods  of 
geological  time  the  sodium  and  potassium  have  been  partly  re- 
moved and  carried  away  to  the  ocean.  Owing  to  the  selective 
action  of  the  colloidal  silicates  remaining  in  the  beds  of  clay  and 
soils  formed  by  the  disintegration  of  the.  rocks,  more  potassium 
than  sodium  has  been  retained,  and  the  element  forms  a  constitu- 
ent of  very  great  importance  in  all  arable  lands.  From  the  soil, 
potassium  is  taken  up  by  all  plants  in  their  growth  and  when 
vegetable  material  is  burned,  the  ash  is  almost  invariably 
alkaline  from  the  presence  in  it  of  potassium  carbonate.  For- 
merly wood  ashes  were  the  most  important  source  of  potassium 
compounds.  The  potassium  carbonate  was  obtained  from  the 
ashes  by  leaching  them  with  water,  and  the  "  lye  "  prepared  in 
this  way  was  used  for  the  domestic  preparation  of  soft  soap. 
The  latter  is  a  concentrated  solution  of  potassium  salts  of  the 
organic  acids  of  ordinary  fats,  such  as  lard  or  tallow.  These 
salts  are  deliquescent  and  cannot  be  readily  brought  to  a  solid 
form,  as  is  the  case  with  the  sodium  salts,  which  form  ordinary 
hard  soap. 

While  beds  of  common  salt  are  found  in  many  different  parts 
of  the  world  and  strong  brines  are  quite  common,  large  deposits 
of  potassium  chloride  (sylvite,  KC1)  and  of  magnesium  potassium 
chloride  (carnallite,  MgCl2.KC1.6.H2O)  have  thus  far  been 

414 


POTASSIUM  415 

found  only  in  Germany,  and  especially  at  Stassfurt.  These 
deposits  now  furnish  the  larger  portion  of  the  potassium  com- 
pounds used  in  the  world,  and  especially  they  furnish  the  potas- 
sium required  to  maintain  the  fertility  of  the  soil  for  raising 
tobacco,  cotton  and  other  crops.  Considerable  amounts .  of 
potassium  compounds  are  found  in  the  seaweeds  of  the  Pacific 
coast,  and  there  is  some  hope  that  they  may  be  profitably  ex- 
tracted from  that  source.  It  is  also  possible  to  prepare  potas- 
sium chloride  on  a  large  scale  from  feldspathic  rocks,  which  are 
abundant  in  some  parts  of  the  United  States. 

Metallic  Potassium  was  prepared  first  by  Sir  Humphry  Davy 
in  1807  at  about  the  same  time  that  he  discovered  metallic  so- 
dium. It  may  also  be  prepared  by  reducing  potassium  carbon- 
ate with  carbon  and  by  the  electrolysis  of  fused  potassium  chlo- 
ride. Potassium  is  a  silver-white  metal  which  tarnishes  instantly 
in  moist  air  and  takes  fire  when  thrown  on  water.  The  hydrogen 
which  is  liberated  burns  with  the  characteristic  violet  flame  of 
potassium.  It  has  a  specific  gravity  of  0.8621  at  20°.  It  melts 
at  62.3°  and  boils  at  about  760°. 

Potassium  Oxide,  K2O,  has  been  prepared  by  the  partial 
oxidation  of  metallic  potassium  in  dry  air  followed  by  distilling 
away  the  excess  of  metal.  It  combines  energetically  with  water 
to  form  potassium  hydroxide,  KOH. 

Potassium  Hydroxide  was  formerly  prepared  by  treating  a 
solution  of  potassium  carbonate  with  slaked  lime : 

K2CO,  +  Ca(OH)2  =  2  KOH  +  CaCO3 

The  reaction  depends,  of  course,  on  the  relative  insolubility 
of  calcium  carbonate. 

Potassium  hydroxide  is  now  prepared  commercially  from 
potassium  chloride  by  electrolysis,  with  the  Castner-Kellner 
and  other  forms  of  apparatus  (p.  402). 

Potassium  hydroxide  is  a  white,  deliquescent  solid.  The  solu- 
tion in  water  has  a  soapy  feel  and  attacks  the  skin  strongly  if 
allowed  to  remain  in  contact  with  it.  The  ordinary  solid  potas- 
sium hydroxide  of  the  laboratory  contains  15  to  20  per  cent  of 


416  A  TEXTBOOK  OF  CHEMISTRY 

water.  Sodium  hydroxide,  on  the  contrary,  is  usually  almost 
anhydrous. 

Solutions  of  potassium  hydroxide  attack  glass  less  than  those 
of  sodium  hydroxide  and  do  not  give  a  precipitate  of  the  carbon- 
ate as  easily  as  the  latter.  For  these  reasons  such  solutions  are 
usually  employed  in  organic  analysis  and  often  in  gas  analysis, 
for  the  absorption  of  carbon  dioxide. 

Potassium  Chloride,  KC1.  The  occurrence  of  potassium  chlo- 
ride and  of  carnal  lite  in  Germany  has  been  mentioned.  It  crys- 
tallizes in  cubes  and  melts  at  about  750°.  It  is  easily  soluble  in 
water.  The  solution  is  often  used  as  a  standard  for  electrical 
conductivity. 

The  crude  salt  is  extensively  employedx  in  fertilizers  and  a 
purer  form  is  used  in  the  manufacture  of  saltpeter,  KNOs. 

Potassium  Chlorate,  KClOs,  may  be  prepared  by  saturating 
a  warm  solution  of  potassium  hydroxide  with  chlorine  (p.  127). 
Practically,  milk  of  lime,  Ca(*OH)2,  is  saturated  with  chlorine, 
forming  calcium  chloride  and  calcium  chlorate,  and  to  this  solu- 
tion potassium  chloride  is  added,  causing  the  separation  of 
potassium  chlorate,  which  is  not  very  easily  soluble.  The 
commercial  reason  for  such  a  procedure  is  apparent.  Potas- 
sium chlorate  is  also  made  by  the  electrolysis  of  a  solution  of 
potassium  chloride  under  such  conditions  that  the  chlorine  and 
potassium  hydroxide  formed  react  with  each  other.  The  final 
result  may  be  expressed  by  the  equation  : 

KC1  +  3  H2O  =  KC1O3  +  3  H2 

Potassium  chlorate  is  used  in  the  preparation  of  oxygen,  in 
medicine  and  in  the  manufacture  of  matches. 

Potassium  Perchlorate,  KC1O4,  is  formed  when  potassium 
chlorate  is  heated  slightly  above  its  melting  point : 

4  KC103  =  3  KC104  +  KC1 

It  is  impossible  to  avoid  some  decomposition  of  the  chlorate 
or  perchlorate  with  evolution  of  oxygen,  but  with  care  a  consid- 
erable portion  of  the  chlorate  may  be  converted  into  the  per- 


POTASSIUM  SALTS  417 

chlorate.  The  latter  is  much  less  soluble  than  the  chloride  or 
chlorate  and  may  be  purified  by  crystallization  from  hot 
water. 

Potassium  Iodide,  KI,  may  be  prepared  by  dissolving  iodine 
in  a  solution  of  potassium  hydroxide.  The  iodate,  KIO3,  formed 
at  the  same  time,  may  be  decomposed  by  heating  the  mixture 
alone,  or,  better,  with  charcoal  or  some  other  reducing  agent. 
The  commercial  salt  often  contains  a  little  iodate,  which  is  very 
objectionable  for  many  laboratory  uses.  The  salt  is  used  in 
medicine  and  for  the  preparation  of  photographic  plates. 

Potassium  Polyiodides.  Solutions  of  potassium  iodide  dis- 
solve iodine  readily,  forming  unstable  polyiodides,  the  one  having 
the  composition  KI3  being  probably  present  in  solutions  of  mod- 
erate concentration.  These  solutions  dissociate  easily  into  potas- 
sium iodide  and  iodine  and  react  in  the  same  manner  as  free 
iodine  toward  reducing  agents.  For  this  reason  such  solutions 
are  used  in  volumetric  analysis.  The  reaction  of  such  a  solution 
with  sodium  thiosulfate  has  been  given  (p.  187). 

Potassium  Sulfate,  K2SO4,  crystallizes  without  water  of  crys- 
tallization. It  melts  at  1080°.  It  forms  double  salts  with 
magnesium  and  calcium,  which  are  found  in  the  potash  deposits 
in  Germany  and  are  an  important  source  of  potassium  com- 
pounds. 

Acid  Potassium  Sulfate,  or  Potassium  Bisulfate,  KHSO4,  can 
be  prepared  by  heating  a  mixture  of  potassium  sulfate  and  sul- 
furic  acid  in  molecular  proportions.  When  heated  gently  it  is 
converted  into  potassium  pyrosulfate,  1^28207,  with  loss  of  water. 
At  a  higher  temperature  the  latter  loses  sulfur  trioxide  and  goes 
back  to  the  normal  sulfate. 

Potassium  pyrosulfate  is  often  used  in  the  laboratory  as  a 
solvent  for  aluminium  oxide,  A^Os,  ferric  oxide,  Fe2Oa,  titanium 
oxide,  TiO2,  and  other  difficultly  soluble  substances.  Sodium 
pyrosulfate  is,  however,  more  suitable  for  this  purpose.  (Hille- 
brand,  Analysis  of  Silicate  and  Carbonate  Rocks,  p.  105.) 

Potassium  Nitrate  or  Saltpeter,  KNO3.  After  the  introduc- 
tion of  gunpowder  into  Europe,  about  1300,  and  especially  after 


418  A  TEXTBOOK  OF  CHEMISTRY 

it  came  into  general  use  in  warfare  in  the  sixteenth  century,  the 
preparation  of  pure  saltpeter  for  use  in  its  manufacture  became 
continually  more  important.  Until  the  nineteenth  century 
saltpeter  was  obtained  almost  exclusively  from  natural  sources, 
where  it  had  been  formed  by  the  decay  of  organic  matter  con- 
taining nitrogen  and  potassium,  in  the  presence  of  nitrifying 
bacteria.  For  a  long  time  considerable  supplies  of  saltpeter  have 
been  obtained  from  India,  where  it  is  formed  in  this  way. 
Calcium  nitrate,  Ca(NOs)2,  which  is  formed  in  a  similar  manner, 
sometimes  occurs  as  an  efflorescence  on  the  walls  of  stables  or  in 
cellars.  By  interaction  with  potassium  carbonate  from  wood 
ashes  saltpeter  is  readily  obtained.  During  the  French  Revolu- 
tion saltpeter  was  often  prepared  in  this  manner.  During  the 
War  of  1812  the  United  States  depended  largely  on  saltpeter 
from  the  Mammoth  Cave,  Kentucky. 

After  the  discovery  of  Chili  saltpeter,  NaNO3,  in  South  Amer- 
ica and  of  potassium  chloride  at  Stassfurt,  the  manufacture  of 
saltpeter  from  these  salts  was  developed.  The  preparation  de- 
pends on  the  fact  that  sodium  chloride  is  about  equally  soluble 
in  hot  or  cold  water,  while  potassium  nitrate  dissolves  in  one 
half  of  its  weight  of  water  at  87°,  but  only  25  parts  of  the  salt  will 
dissolve  in  100  parts  of  water  at  15°.  If  potassium  chloride  is 
added  in  molecular  proportions  to  a  concentrated,  hot  solution  of 
sodium  nitrate,  it  will  pass  into  solution,  and  sodium  chloride, 
the  least  soluble  of  the  four  salts  (NaNO3,  KC1,  NaCl,  KNO3) 
present,  will  separate.  From  the  mother  liquors  potassium  ni- 
trate will  separate  on  cooling,  since  that  is  the  least  soluble  con- 
stituent in  the  cool  solution.  For  the  manufacture  of  gunpowder 
a  salt  entirely  free  from  chloride  must  be  prepared  by  recrys- 
tallization  and  centrifugal  drainage. 

Potassium  nitrate  crystallizes  in  rhombic  prisms.  It  melts  at 
339°.  It  is  used  in  the  manufacture  of  gunpowder  and  in  the 
curing  of  meats,  especially  of  salt  beef.  It  imparts  to  the  meat 
a  desirable  reddish  color.  Taken  in  considerable  quantities  it 
is  a  poison. 

Gunpowder  is  a  mixture  of  about  75  parts  of  saltpeter,  13  parts 


POTASSIUM  SALTS  419 

of  charcoal  and   12  parts  of  sulfur.     This   corresponds   very 
nearly  to  the  equation  : 

2  KN03  +  3  C  -f  S  =  K2S  +  N2  +  3  CO2 

The  explosion  depends  on  the  fact  that  the  oxygen  for  burning 
the  carbon  is  contained  in  the  mixture  and  on  the  large  volume  of 
the  nitrogen  and  carbon  dioxide  formed  in  the  reaction.  The 
heat  of  combustion  also  raises  these  gases  to  a  high  temperature, 
increasing  the  force  of  the  explosion.  In  the  burning  of  the 
gunpowder  the  grains  burn  from  the  surface  inward,  and  the 
speed  of  the  combustion  is  closely  connected  with  the  size  of 
the  grains.  For  use  in  large  ordnance,  hexagonal  blocks  an  inch 
or  more  in  diameter  are  used  to  secure  slower  combustion  and 
allow  time  for  the  heavy  shot  to  gain  momentum  before  the  full 
force  of  the  explosive  is  developed.  The  failure  of  guncotton 
when  it  was  first  tried  in  firearms  was  partly  due  to  the  too  rapid 
burning  of  the  material,  which  caused  the  guns  in  which  it  was 
used  to  burst.  This  difficulty  was  finally  overcome  by  giving 
the  "  smokeless  powder  "  made  from  guncotton  a  dense  form, 
somewhat  resembling  that  of  ordinary  gunpowder. 

Potassium  Nitrite,  KNO2,  is  prepared  by  the  reduction  of  the 
nitrate  with  lead,  iron  or  sometimes  with  charcoal  or  sulfur.  It 
is  very  easily  soluble  in  water  and  is  used  in  the  laboratory  as  a 
reagent  for  cobalt,  with  which  it  forms  a  difficultly  soluble  com- 
plex salt,  Co(N02)3.3  KN02  or  K3Co(NO2)6.  The  forma- 
tion of  the  same  salt  or  of  a  similar  salt  containing  silver, 
K2AgCo(NO2)6,  may  also  be  used  as  a  test  for  detecting 
potassium. 

Potassium  Carbonate,  K2CO3,  was  formerly  obtained  by  leach- 
ing wood  ashes.  The  salt  is  also  obtained  in  the  scouring  of  wool 
and  from  the  residual  sirups  of  the  beet  sugar  manufacture  after 
the  sugar  of  the  sirups  has  been  converted  into  alcohol.  It  is 
also  made  from  potassium  chloride  by  processes  similar  to  those 
used  in  manufacturing  sodium  carbonate. 

Potassium  carbonate  is  a  deliquescent  salt,  differing  in  this 
respect  very  markedly  from  sodium  carbonate.  The  anhydrous 


420  A  TEXTBOOK  OF  CHEMISTRY 

salt  melts  at  about  890°.  It  is  used  in  making  soft  soap  and  hard 
glass. 

Potassium  Bicarbonate  or  Saleratus,  KHCOs,  is  easily  pre- 
pared by  passing  carbon  dioxide  into  a  concentrated  solution 
of  potassium  carbonate.  It  was  formerly  used  in  cooking,  but 
has  been  entirely  displaced  by  the  cheaper  and  more  suitable 
sodium  bicarbonate,  NaHCO3.  Potassium  bicarbonate  dis- 
solves in  about  three  parts  of  water,  being  much  more  easily 
soluble  than  the  sodium  salt.  The  solution  is  nearly  neutral 
to  phenolphthalein,  but  loses  carbon  dioxide  and  becomes  alka- 
line on  boiling. 

Potassium  Cyanide,  KCN,  can  be  prepared  by  heating  potas- 
sium ferrocyanide : 

K4FeC6N6  =  4  KCN  +  Fe  +  2  C  +  &* 

By  heating  the  salt  with  metallic  sodium  a  mixture  of  potas- 
sium and  sodium  cyanides  is  obtained  and  all  of  the  cyanogen 
of  the  original  salt  can  be  saved.  This  mixture  may  be  used 
for  nearly  all  purposes  to  which  potassium  cyanide  is  applied  and 
especially  for  the  extraction  of  gold  from  its  ores. 

Ammonium,  NH4,  is  a  group  which  so  closely  resembles  potas- 
sium in  the  salts  which  it  forms  with  acid  radicals  that  it  seems 
desirable  to  speak  of  these  salts  at  this  point.  It  has  not  been 
found  possible  to  separate  ammonium,  NH4,  by  itself,  but  if  a 
solution  of  ammonium  chloride  is  poured  on  some  sodium 
amalgam  the  reaction  represented  by  the  equation  : 

NH4C1  +  Na(Hg)  =  NaCl  +  NH4(Hg) 

takes  place  and  it  can  be  shown  by  the  electrical  properties 
and  by  the  effect  in  reducing  metals  from  their  salts  that  a  small 
amount  of  ammonium  amalgam  is  formed.  The  substance  is, 
however,  extremely  unstable  and  decomposes  rapidly  into  am- 
monia, NHs,  hydrogen  and  mercury. 

Ammonium  Hydroxide,  NH4OH.  Solutions  of  ammonia  in 
water  seem  to  involve  the  following  equilibria : 

NH3  +  H20  ^±  NH4OH  ;±  NH4+  +  GET 


AMMONIUM    SALTS  421 

The  ratio  between  these  various  substances  has  not  been  deter- 
mined with  any  degree  of  certainty.  It  is  known,  however,  that 
the  concentration  of  the  hydroxide  ions  is  small  in  comparison 
with  the  concentration  in  a  solution  of  sodium  or  potassium 
hydroxide  of  equivalent  molecular  concentration.  In  other 
words,  ammonium  hydroxide  is  a  comparatively  weak  base. 

Ammonium  Chloride,  NH4C1,  is  prepared  by  neutralizing  the 
aqueous  gas  liquors,  obtained  in  the  manufacture  of  coal  gas  or 
coke,  with  hydrochloric  acid.  It  is  partially  purified  by  subli- 
mation. When  heated  to  about  350°  under  ordinary  conditions, 
it  is  converted  into  a  gas  which  consists  of  a  mixture  of  ammonia 
and  hydrochloric  acid,  the  weight  of  a  gram  molecular  volume 
of  the  gas  being  only  about  26.8  grams  instead  of  53.5  grams, 
as  would  be  expected  from  the  formula.  This  abnormal  density 
was  at  one  time  used  as  an  argument  against  Avogadro's  hy- 
pothesis. It  has  been  shown  by  diffusion  experiments,  however, 
that  the  gas  is  a  mixture,  as  the  ammonia,  which  is  the  lighter 
of  the  two  constituents,  diffuses  away  more  quickly  than  the 
hydrochloric  acid.  Finally,  many  years  after  the  theory  of  dis- 
sociation had  been  universally  accepted  as  the  correct  explana- 
tion of  the  abnormal  density,  it  was  shown  that  very  carefully 
dried  ammonium  chloride  may  be  converted  into  a  vapor  without 
dissociation,  and  that  a  gram  molecular  volume  of  this  vapor 
weighs  about  53.5  grams. 

Ammonium  Sulfide,  (NH^S,  is  prepared  by  passing  hydrogen 
sulfide,  H2S,  into  a  solution  of  ammonia.  The  most  convenient 
method  is  to  take  a  known  quantity  of  a  10  per  cent  solution  of 
ammonia  (sp.  gr.  0.96)  and  pass  into  it  the  hydrogen  sulfide 
generated  by  the  action  of  one  fifth  of  its  volume  of  concentrated 
sulfuric  acid  upon  an  excess  of  ferrous  sulfide,  FeS,  contained  in  a 
bottle  holding  ten  volumes  of  water  for  one  volume  of  the  acid. 
The  gas  should  be  generated  rapidly  and  well  washed  with  water 
before  entering  the  ammonia. 

Ammonium  Hydrosulfide,  NH4SH.  If  hydrogen  sulfide  in 
excess  is  passed  into  a  solution  of  ammonia,  the  hydrosulfide  is 
formed.  When  a  solution  of  either  ammonium  sulfide  or  of 


422  A  TEXTBOOK  OF  CHEMISTRY 

the  hydrosulfide  is  exposed  to  the  air,  especially  if  exposed  also 
to  light,  the  sulfides  are  oxidized,  sulfur  separates  and  am- 
monium hydroxide  is  regenerated.  The  principal  action  is 
parallel  to  that  of  air  on  hydrogen  sulfide  water : 

H2S  +  O  =  H2O  +  S 
(NH4)2S  +  O  +  H2O  =  2  NH4OH  +  S 

The  sulfur  liberated  in  this  manner  will,  for  a  time,  dis- 
solve in  the  ammonium  sulfide,  forming  poly  sulfides,  (NH4)2S2, 
(NH^A,  etc.  These  polysulfides  give  to  the  solution  a  yellow 
color.  After  the  oxidation  has  gone  beyond  a  certain  point,  the 
separated  sulfur  no  longer  finds  any  ammonium  sulfide  with 
which  to  combine  and  begins  to  separate  in  the  free  state.  The 
solution  is  then  no  longer  fit  for  use  as  a  reagent. 

Ammonium  sulfide  is  used  in  the  laboratory  to  precipitate 
those  metals  whose  sulfides  are  too  soluble  for  precipitation  in 
acid  solution  but  sufficiently  insoluble  for  precipitation  in  the 
presence  of  a  base.  The  polysulfide  is  used  to  convert  stannous 
sulfide,  SnS,  into  stannic  sulfide,  SnS2,  and  to  dissolve  the  latter 
in  separating  it  from  lead,  bismuth  and  other  metals.  Either 
may  be  used  to  dissolve  arsenious  sulfide,  As2S3,  or  antimony 
sulfide,  Sb2S3  (p.  261). 

Ammonium  Sulfate,  (NH4)2SO4,  is  prepared  in  a  crude  form 
by  neutralizing  the  ammoniacal  liquors  of  the  gas  works  with 
sulfuric  acid.  As  sulfuric  acid  is  the  cheapest  of  the  commercial 
acids,  this  salt  is  often  prepared  to  put  the  ammonia  into  suit- 
able form  for  transportation  or  for  use  in  fertilizers.  Ammonia 
can,  of  course,  be  readily  regenerated  from  it  by  treatment  with 
lime,  CaO. 

Ammonium  Nitrate,  NH4NO3,  may  be  prepared  by  neutraliz- 
ing nitric  acid  with  ammonia  or  ammonium  carbonate  and 
evaporating  the  solution  to  crystallization.  It  is  easily  soluble 
m  water.  At  166°  it  melts  and  decomposes  into  nitrous  oxide, 
N2O,  and  water.  The  decomposition  is  exothermic,  and  if  the 


AMMONIUM  SALTS  423 

. 

temperature  is  too  high  or  if  a  large  amount  of  the  salt  is  heated 

at  once,  the  reaction  may  become  explosive  : 

NH4N03  =  N2O  +  2  H20  +  29,500  cal. 


The  salt  is  also  used  as  a  very  important  constituent  of  modern 
explosives. 

Ammonium  Nitrite,  NH4NO2,  is  a  deliquescent,  very  unstable 
salt,  which  decomposes  easily  into  nitrogen  and  water. 

Ammonium  Sodium  Hydrogen  Phosphate,  or  microcosmic  salt, 
NaNH4HPO4.4H2O,  is  used  in  blowpipe  analysis  to  furnish  a 
bead  of  sodium  metaphosphate,  NaPO3,  which,  when  hot,  will 
dissolve  many  metallic  oxides,  giving  characteristic  colors.  The 
bead  does  not  dissolve  silica,  SiO2. 

Ammonium  Carbonate,  (NH4)2CO3.  A  salt  known  commer- 
cially as  "  ammonium  carbonate  ""  is  prepared  by  heating  a 
mixture  of  calcium  carbonate  and  ammonium  sulfate.  It  con- 
sists of  a  mixture  of  ammonium  bicarbonate,  NH4HCO3,  and 

ammonium  carbamate,  NH4  —  O  —  C—  NH2.  The  second  of 
these  salts  may  be  prepared,  also,  by  the  direct  union  of  carbon 
dioxide  and  ammonia.  By  dissolving  the  commercial  carbonate 
in  water  and  adding  ammonia,  the  bicarbonate,  NH4HCO3,  is 
changed  to  the  normal  carbonate,  (NH4)2CO3.  The  carbamate 
is  also  soon  hydrolyzed  to  the  normal  carbonate  : 

NH4—  O—  C—  NH2+H20=  NH4—  O—  C^O—  NH4 
A  solution  prepared  in  this  way  is  used  as  the  ordinary  labora- 

tory reagent. 

Ammonium  Bicarbonate,  NH4HCO3,  is  prepared  by  passing 

carbon  dioxide  into  a  solution  of  ammonia  as  one  of  the  opera- 

tions of  the  ammonia-soda  process  (p.  412). 
Ammonium  Chloroplatinate,  (NH4)2PtCl6,  is  a  very  difficultly 

soluble  salt  closely  resembling  the  corresponding  potassium  salt. 

Similar  compounds  are  formed  from  many  amines,  compounds 

in  which  one  or  more  hydrogen  atoms  of  ammonia  have  been  re- 

placed by  organic  radicals. 


424  A  TEXTBOOK  OF   CHEMISTRY 

Rubidium,  Rb,  85.45,  and  Caesium,  Cs,  132.81.  In  1860  and 
1861  Bunsen  and  Kirchoff  in  applying  their  newly  discovered 
method  of  spectrum  analysis  to  the  study  of  a  mineral  water 
from  Durkheim  found  some  spectral  lines  which  did  not  corre- 
spond to  those  of  any  known  element.  In  order  to  obtain  enough 
material  to  study  the  compounds  of  the  new  elements,  40  tons 
of  the  water  were  evaporated  and  the  compounds  of  rubidium 
and  caesium  were  extracted  from  the  residues.  Rubidium  was 
named  from  the  Latin  word  rubidus,  meaning  red,  and  caesium 
from  the  Latin  ccesiw,  the  blue  of  the  sky,  because  of  the  red 
lines  in  the  spectrum  of  the  former  and  the  blue  lines  in  the 
spectrum  of  the  latter.  Rubidium  is  found  to  the  amount  of 
about  0.025  per  cent  in  carnallite  (KMgCl3.6H2O)  and  as  a 
million  and  a  half  tons  of  this  mineral  are  worked  over  annually 
for  the  preparation  of  potassium  compounds  it  would  be  possible 
to  obtain  very  considerable  quantities  of  the  element. 

Metallic  rubidium  melts  at  38.5°,  caesium  at  26.5°,  the  lowest 
melting  point  of  any  metal  except  mercury.  Their  compounds 
resemble  those  of  potassium  in  their  general  properties.  The 
chloroplatinates,  Rb2PtCl6,  and  Cs2PtCl6,  and  the  alums, 
RbAl(SO4)2.12  H2O  and  CsAl(SO4)2.12  H2O,  are  less  soluble 
than  the  corresponding  potassium  compounds  and  are  used  in  sep- 
arating the  elements.  A  chloroiodide  of  caesium,  CsCl2I,  is  also 
especially  useful  for  the  preparation  of  pure  caesium  compounds. 

Spectrum  Analysis.  Early  in  the  nineteenth  century  Fraun- 
hofer  pointed  out  that  when  a  solar  spectrum  is  produced  in  such 
a  manner  that  the  colors  are  sharply  separated  from  each  other, 
the  spectrum  is  crossed  by  a  series  of  dark  lines.  During  the 
years  from  1820  to  1860  several  different  observers  noticed  the 
characteristic  colors  imparted  to  flames  by  different  elements  and 
the  spectra  of  bright  lines  given  by  these  colored  flames  and  also 
those  given  by  metals  which  are  vaporized  by  the  electric  spark. 
It  was  not  till  1860,  however,  that  all  of  these  phenomena  were 
brought  into  clear  relationship  by  the  classical  researches  of 
Kirchoff  and  Bunsen,  which  culminated  in  the  discovery  of 
rubidium  and  caesium. 


THE  SPECTROSCOPE 


425 


The  simplest  form  of  spectroscope  is  shown  in  Fig.  97.  The 
prism  A  is  of  glass  having  a  high  dispersive  power.  A  narrow 
slit  at  B  is  illuminated  by  the  flame  or  light  which  is  to  be  exam- 
ined. Between  the  slit  and  the  prism  is  placed  a  lens  at  (7, 
which  renders  the  rays  of  light  from  the  slit  parallel  before  they 
B 


Fig.  97 

reach  the  prism.  The  prism  is  set  in  such  a  manner  that  the 
angle  of  incidence  on  the  first  face  is  the  same  as  the  angle  of 
emergence  from  the  second  face,  as  this  gives  the  purest  spectrum. 
The  light  is  examined  by  means  of  the  telescope  D.  A  scale 
placed  at  E,  illuminated  by  a  light  placed  before  it  and  whose 
image  is  reflected  from  the  surface  of  the  prism,  serves  to  locate 
the  position  of  the  lines. 

Such  a  flame  as  that  of  acetylene  gives  a  continuous  spec- 
trum, indicating  that  molecules  of  solid  carbon  in  the  white 
flame  are  vibrating  in  all  possible  periods  required  to  give  white 
light.  If  a  Bunsen  flame  is  placed  before  the  slit  and  some  com- 
pound of  sodium,  as  sodium  chloride,  is  introduced,  the  flame 
assumes  a  brilliant  yellow  color,  and  with  a  single  prism  spec- 
troscope the  spectrum  consists  of  a  single,  bright  yellow  line. 


426  A  TEXTBOOK  OF  CHEMISTRY 

A  spectroscope  having  several  prisms,  or  a  spectroscope  using 
a  metallic  mirror  ("  grating  ")  ruled  with  many  thousands  of 
equidistant  lines  and  which  gives  a  diffraction  spectrum,  will 
separate  the  line  into  two  lines  situated  close  together.  The 
wave  lengths  of  the  lines  are  0.5896  and  0.5890  microns,  the 
micron  being  the  thousandth  part  of  a  millimeter.  The  physi- 
cal significance  of  these  lines  seems  to  be  that  under  the  con- 
ditions of  the  flame  either  the  sodium  atoms  as  a  whole  or, 
more  probably,  portions  of  the  sodium  atoms  or  electrons  within 
or  around  them  vibrate  at  a  definite  rate,  which  is  independent 
of  the  temperature.  This  rate  is  almost  inconceivably  rapid. 
The  velocity  of  light  is  about  300,000  kilometers  per  second. 
This  is  equal  to  3  X  1014  microns,  and  since  the  wave  length  of 
the  sodium  light  is  only  0.59  micron,  the  number  of  vibrations 

o   vx    1Q14 

per   second    must   be  approximately  =  5  X  1014  per 

u.oy 

second. 

There  are  two  dark  lines  in  the  solar  spectrum  which  coincide 
exactly  with  the  bright  yellow  lines  of  the  sodium  spectrum. 
This  is  explained  by  supposing  the  interior  of  the  sun  to  be  an 
incandescent  mass  which  gives  out  light  vibrations  of  all  wave 
lengths  corresponding  to  the  visible  spectrum.  The  photo- 
sphere of  the  sun,  on  the  other  hand,  consists  of  a  gaseous  en- 
velope or  atmosphere  containing  many  different  elements,  among 
these  sodium.  The  sodium  atoms,  if  they  have  the  power  of 
producing  light  waves  in  the  ether  by  their  vibrations,  must  also 
be  able  to  absorb  waves  of  the  same  length  from  the  ether, 
exactly  as  a  tuning  fork  is  set  in  vibration  by  sound  waves  of  its 
own  pitch,  while  waves  of  a  different  pitch  do  not  affect  it.  The 
sodium  atoms  in  the  photosphere,  therefore,  absorb  the  waves 
of  their  own  particular  rate ;  and  while  they  give  the  energy  ab- 
sorbed back  again  to  the  ether,  they  dissipate  the  energy  by 
spreading  it  in  all  directions  instead  of  allowing  it  to  pass  on 
toward  the  observer.  The  result  is  that  the  portion  of  the 
spectrum  corresponding  to  the  sodium  vibrations  will  be  rela- 
tively dark.  By  means  of  this  principle  it  has  been  possible  to 


THE  SPECTROSCOPE  427 


show  that  more  than  thirty  elements  found  on  the  earth  are  found 
also  in  the  sun.  One  of  these  elements  (helium),  indeed,  was 
discovered  in  the  sun  before  it  was  found  on  the  earth. 

For  the  purpose  of  comparing  spectra  it  is  convenient  to  place 
a  right-angled  prism  before  the  slit  of  the  spectroscope  in  such  a 
manner  as  to  cover  one  half  of  it.  This  may  be  made  to  reflect 
the  light  from  a  second  flame  into  the  slit  in  such  a  way  that 
the  spectrum  from  one  flame  will  occupy  the  upper  half  of  the 
field  of  vision  while  the  spectrum  from  the  other  flame  will  occupy 
the  lower  half.  In  this  manner  the  coincidence  of  lines  in  the 
two  spectra  may  be  readily  observed. 

In  another  form,  known  as  the  direct  vision  spectroscope,  a 
series  of  prisms  of  different  kinds  of  glass  are  so  combined  that 
one  kind  of  glass  counterbalances  the  mean  refractive  index  of  the 
other,  while  the  dispersive  effects .  are  not  counterbalanced. 
The  effect  is  exactly  the  reverse  of  that  in  an  achromatic  lens. 
Such  spectroscopes  are  especially  suitable  for  the  detection  of  the 
alkali  and  alkali-earth  metals  in  qualitative  analysis.  By  means 
of  the  spectroscope  it  is  possible  to  detect  3000000  milligram 
of  soqjium.  Only  the  methods  used  in  studying  radioactive  sub- 
stances are  more  sensitive  than  this. 

To  obtain  the  spectra  of  iron,  copper  and  other  metals,  which 
are  volatile  only  at  high  temperatures,  electric  sparks  from  a 
Rumkhorf  coil  are  passed  between  terminals  of  the  metal,  or, 
in  some  cases,  between  platinum  wires,  one  of  which  is  in  a  small 
cup  containing  a  solution  of  a  salt  of  the  metal.  The  spectra 
of  gases  are  observed  in  Pliicker  tubes,  which  have  a  narrow 
portion  through  which  the  electric  discharge  is  passed. 


CHAPTER  XXV 

THE  ALTERNATE  METALS   OF  GROUP  I.   COPPER, 
SILVER,   GOLD.    PHOTOGRAPHY 

COPPER,  silver  and  gold,  which  alternate  with  potassium,  rubid- 
ium and  caesium  in  Group  I  of  the  Periodic  System,  are  in 
almost  the  greatest  possible  contrast  with  those  metals.  The 
alkali  metals  are  light.  They  melt  and  volatilize  at  compara- 
tively low  temperatures  and  they  react  violently  with  water  at 
ordinary  temperatures.  Copper,  silver  and  gold  are  heavy 
metals,  all  three  melt  between  960°  and  1083°  and  none  of  them 
decomposes  water,  even  at  high  temperatures.  They  are  also 
the  best  conductors  of  electricity  that  we  have.  As  they  do 
not  decompose  water,  all  three  of  these  metals  are  found  free 
in  nature.  All  three  have  been  known  and  used  since  very 
early  times. 

Copper,  Cu,  63.57.  Occurrence.  Copper  is  found  free  in 
nature,  especially  in  the  Lake  Superior  region. 

It  is  found  as  copper  pyrites,  or  chalcopyrite,  CuFeS2,  a  mineral 
closely  resembling  iron  pyrites  in  superficial  appearance  but 
having  a  different  crystalline  form  and  usually  showing  blue, 
red  or  green  colors,  owing  to  superficial  changes.  Bornite, 
CuaFeSa,  chalcocite,  cuprous  sulfide,  Cu2S,  and  malachite, 

/O— Cu— OH 
CuCO3.Cu(OH)2,    or    C<X  » a  Dasic  carbonate,  are 

\O-Cu— OH 
the  other  most  important  minerals  containing  copper. 

Metallurgy.  The  most  common  ores  of  copper  contain  either 
the  sulfide,  Cu2S,  or  copper  pyrites,  CuFeS2.  If  the  ore  is 
poor  in  copper,  it  is  sometimes  concentrated  by  crushing  it 
and  washing  away  a  part  of  the  lighter  minerals  with  water. 
The  concentrated  ore  is  then  roasted  in  a  furnace  with  the 

428 


COPPER  429 

addition  of  sand,  if  enough  silica  is  not  already  present.  The 
iron  is  partly  oxidized  to  ferrous  oxide,  FeO,  which  combines 
with  the  silica,  SiO2,  to  form  a  fusible  silicate  or  slag,  Fe2SiO4. 
The  cuprous  sulfide  and  some  of  the  ferrous  sulfide  melt  and 
sink  to  the  bottom  of  the  furnace  beneath  the  slag,  which  is 
much  lighter.  The  mixture  of  sulfides  obtained  in  this  way  is 
drawn  off  and  is  known  as  copper  matte.  Similar  operations  are 
now  often  carried  out  in  a  blast  furnace  somewhat  similar  to 
that  used  in  the  manufacture  of  pig  iron  (p.  543).  In  the  older 
processes  the  ore  was  usually  carried  through  a  long  series  of 
complicated  operations  for  the  purpose  of  securing  a  compara- 
tively pure  matte  and  reducing  the  latter  to  metallic  copper. 
In  the  United  States  these  processes  have  been  very  much 
shortened  by  the  Use  of  a  modification  of  the  Bessemer  converter 
(p.  547)  for  the  reduction  of  the  matte.  The  molten  matte 
from  the  roasting  furnace  is  poured  directly  into  the  converter, 
where  it  is  subjected  to  a  blast  of  air  mixed  with  fine  sand  or 
silica.  The  sulfur  of  the  ferrous  sulfide  and  cuprous  sulfide  is 
burned  out,  the  heat  of  the  combustion  maintaining  the  tem- 
perature of  the  converter.  The  ferrous  oxide  combines  with 
the  sand  to  form  a  slag  of  ferrous  silicate,  while  the  copper 
melts  and  may  be  cast  into  plates.  The  principal  reactions 
may  be  expressed  as  follows  : 

f    2FeS  +  3O2  =  2FeO  +  2SO2 
1st  btage  |  2FeO+  SiO2  =  Fe2SiO4 

r     Cu2S  +  2O2  =  2CuO  +  SO2 
2d  Stage  1 2  CuO  +  Cu2S  =  4  Cu  +  SO2 
I  or  Cu2S  +  O2  =  2  Cu  +  SO2 

Electrolytic  Refining  of  Copper.  The  copper  obtained  by 
means  of  the  Bessemer  converter  or  by  any  of  the  more  com- 
plicated furnace  methods  usually  contains  a  small  amount  of 
gold  and  silver  and  larger  quantities  of  arsenic,  lead  and  other 
metals  which  render  it  unfit  for  most  industrial  uses  and  es- 
pecially for  use  in  electrical  conductors.  Nearly  all  of  the 
copper  is  now  refined  electrolytically.  The  plates  of  crude 


430 


A  TEXTBOOK  OF  CHEMISTRY 


copper  are  suspended,  upright,  in  a  long  tank  (Fig.  98)  filled 
with  a  solution  of  copper  sulfate.  By  connecting  the  two  end 
plates  with  the  poles  of  a  dynamo  the  current  flowing  through 
the  system  will  cause  each  plate  to  become  negative  on  one  side 
and  positive  on  the  other.  On  the  positive  side  the  copper  will 


Fig.  98 

pass  into  solution  as  cupric  ion,  Cu++.  Gold,  silver,  bismuth 
and  some  other  impurities  in  the  copper  fail  to  dissolve  and  are 
collected  as  "  slimes  "  and  treated  for  the  recovery  of  silver  and 
gold.  On  the  negative  side  of  each  plate  nearly  pure  copper 
will  be  deposited.  When  all  of  the  original  copper  has  been 
dissolved  the  plates  of  pure  copper  are  removed  and  new  plates 
of  crude  copper  put  in  their  place. 

The  production  of  copper  in  the  United  States  in  1910  was 
500,000  tons,  worth  $137,000,000. 

Properties  of  Copper.  Copper  is  red  as  ordinarily  seen  by 
reflected  light.  In  very  thin  films  it  transmits  green  light.  It 
melts  at  1083°  (1063°  in  air,  because  of  oxidation),  boils  at  2310° 
and  has  a  density  of  8.93.  It  is  the  best  conductor  of  electricity 
of  the  cheaper  metals.  Its  conductance  is  very  much  impaired, 
however,  by  the  presence  of  small  amounts  of  other  metals. 
Arsenic  is  especially  harmful,  0.03  per  cent  lowering  the  con- 
ductance by  14  per  cent.  A  table  of  conductances  for  the  more 
common  metals  will  be  found  at  the  close  of  this  chapter. 

When  exposed  to  the  weather,  copper  slowly  covers  itself 


ALLOYS   OF  COPPER  431 

with  a  green  coating  of  basic  carbonate  of  the  composition  of 
malachite,  CuCO3.Cu(OH)2.  When  heated  in  the  air  a  coating 
of  the  black  oxide,  CuO,  is  formed,  which  comes  off  in  scales 
on  cooling.  In  the  absence  of  air,  copper  is  not  affected  by 
hydrochloric  or  dilute  sulfuric  acid.  It  dissolves  in  nitric  acid 
as  copper  nitrate,  Cu(NO3)2,  with  evolution  of  nitric  oxide,  NO, 
nitrous  oxide,  N2O,  nitrogen  peroxide,  NO2,  or  a  mixture  of 
these  according  to  the  concentration  of  the  acid  and  the  tem- 
perature. Hot,  concentrated  sulfuric  acid  dissolves  it  as  cupric 
sulfate,  CuSO4,  while  sulfur  dioxide  is  evolved. 

Copper  is  used  for  electrical  conductors,  for  the  sheathing  of 
ships  and  for  much  of  the  apparatus  used  in  the  fermentation 
industries. 

Alloys  of  Copper.  Copper  is  used  in  a  great  variety  of  alloys, 
the  two  most  important  being  brass  and  bronze.  Brass  con- 
tains 60  to  70  per  cent  of  copper  and  40  to  30  per  cent  of  zinc, 
with  usually  small  amounts  of  lead,  tin  and  iron.  Bronze  is 
primarily  an  alloy  of  copper  and  tin,  usually  with  80  to  90  per 
cent  of  copper ;  it  contains,  in  most  cases,  small  amounts  of  zinc 
and  lead.  Bronze  is  used  for  bells  and  statuary  and  was  for- 
merly also  used  for  cannon.  Phosphor  bronze  contains  from 
0.25  to  2.5  per  cent  of  phosphorus,  which  makes  it  hard  and 
suitable  for  bearings  in  machinery.  It  has  recently  been  dis- 
covered that  the  addition  of  a  very  small  amount  of  copper  to 
iron  or  steel  greatly  increases  its  resistance  to  corrosion. 

Copper  Hydroxide,  Cu(OH)2,  is  precipitated  as  a  blue,  amor- 
phous compound  when  a  solution  of  sodium  hydroxide  is  added 
to  a  solution  of  copper  sulfate  or  any  similar  salt.  If  the  solu- 
tion is  heated  to  boiling,  the  hydroxide  is  decomposed  into  the 
black  cupric  oxide  and  water : 

Cu(OH)2  =  CuO  +  H2O 

The  instability  of  copper  hydroxide  is  in  very  marked  con- 
trast with  the  stability  of  the  hydroxides  of  the  alkali  metals, 
but  the  hydroxides  of  silver  and  mercury  are  still  less  stable 
and  have  never  been  prepared  as  pure  compounds. 


432  A  TEXTBOOK  OF  CHEMISTRY 

Cupric  Oxide,  CuO,  may  be  prepared,  as  just  described,  by 
the  decomposition  of  the  hydroxide,  or  by  the  decomposition 
of  the  nitrate,  but  neither  method  will  give  a  pure  compound. 
The  copper  oxide  which  is  used  for  organic  analysis  is  usually 
prepared  by  calcining  copper  wire  in  the  air  for  a  long  time. 
Copper  oxide  ..obtained  by  heating  the  nitrate  retains  nitrogen 
obstinately  and  is  not  suitable  for  use  in  the  determination  of 
nitrogen  in  organic  compounds. 

Cuprous  Oxide,  Cu2O,  is  formed  as  a  bright  red  precipitate 
by  the  action  of  reducing  agents  and  especially  by  the  action 
of  hydroxylamine,  hydrazine,  glucose,  fructose  or  lactose  on  a 
solution  of  copper  sulfate,  sodium  hydroxide  and  potassium 
sodium  tartrate  (p.  335). 

Cupric  Chloride,  CuCl2.2H2O,  is  a  bright  green  salt  obtained 
by  dissolving  cupric  oxide  in  hydrochloric  acid.  The  double 
salt,  CuCl2.2KC1.2  H2O,  is  used  to  dissolve  iron  or  steel  without 
evolution  of  gas  for  the  determination  of  carbon. 

Copper  chloride  dissolves  in  a  small  amount  of  water  to  a 
green  solution  similar  in  color  to  the  crystallized  salt.  On 
dilution  the  color  changes  to  the  blue  color  characteristic  of  all 
solutions  containing  the  cupric  ion,  Cu++.  If  not  too  dilute, 
concentrated  hydrochloric  acid  will  restore  the  green  color, 
because  the  excess  of  chloride  ions  causes  a  decrease  in  the 
ionization  of  the  cupric  chloride. 

Cuprous  Chloride,  Cu2Cl2,  may  be  easily  prepared  by  digesting 
a  solution  of  cupric  chloride,  CuCl2,  in  concentrated  hydro- 
chloric acid  with  copper  turnings : 

CuCl2  +  Cu  =  Cu2Cl2 

Cuprous  chloride  is  very  difficultly  soluble  in  water  but  dis- 
solves rather  easily  in  concentrated  hydrochloric  acid,  owing  to 
the  formation  of  chlorocuprous  acid,  HCu2Cl3.  On  dilution 
this  is  dissociated  into  its  components  and  cuprous  chloride 
is  precipitated.  The  insolubility  of  cuprous  chloride,  Cu2Cl2,  is 
parallel  to  that  of  silver  chloride,  AgCl  [Ag2Cl2],  and  mercurous 
chloride,  Hg2Cl2. 


COPPER  SALTS  433 

Cuprous  chloride  dissolves  easily  in  ammonia,  also.  It  is 
supposed  that  a  complex  molecule  of  the  form,  [Cu(NH3)n]Cl, 
is  formed,  but  the  exact  composition  of  this  molecule  has  not 
been  determined. 

Solutions  of  cuprous  chloride  either  in  hydrochloric  acid  or 
in  ammonia  absorb  carbon  monoxide  and  are  used  for  that 
purpose  in  gas  analysis. 

Sodium  hydroxide  decomposes  cuprous  chloride  with  the  for- 
mation of  cuprous  oxide.  Cuprous  hydroxide,  CuOH,  has  not 
been  prepared  as  a  definite  compound,  and  there  is  some  reason 
to  suppose  that  it  could  exist  only  at  low  temperatures. 

*  Cuprous  Iodide,  Cu2I2.  If  potassium  iodide  is  added  to  an 
acid  solution  of  a  cupric  salt,  cuprous  iodide  is  precipitated  and 
iodine  is  liberated : 

2  Cu(NO3)2  +  4  KI  =  Cu2I2  +  4  KNO3  +  I2 
The  reaction  is  quantitative,  and  by  titrating  the  iodine  with 
a  solution  of  sodium  thiosulfate  the  amount  of  copper  present 
may  be  determined.     Cuprous  iodide  is  white  and  nearly  in- 
soluble in  water. 

Cupric  Sulfide,  CuS,  is  formed  as  a 'black  precipitate  on  pass- 
ing hydrogen  sulfide  into  an  acid  solution  of  a  cupric  salt.  It 
is  very  insoluble  in  water  or  dilute  acids,  but  dissolves  readily 
in  warm  nitric  acid,  the  sulfur  separating  mostly  in  the  free 
state.  On  ignition  in  a  current  of  hydrogen  it  is  converted  into 
cuprous  sulfide,  Cu2S,  which  has  the  same  composition  as  the 
mineral  chalcocite. 

Copper  Sulfate,  or  Blue  Vitriol,  CuSO4.5H2O,  forms  deep  blue 
crystals  of  the  triclinic  system.  At  120°-140°  these  lose  four 
fifths  of  their  water,  leaving  the  hydrate,  CuSO4.H2O.  This 

/°\    ^° 

has,  very  probably,  the  structure  Cu<      )>S^-OH.     At  240° 

XX      \OH 

the  last  molecule  of  water  may  be  expelled,  leaving  the  anhydrous 
salt,  CuSO4,  as  a  white  powder.  Copper  sulfate  is  the  most 
common  salt  of  copper.  It  is  used  in  the  electrolytic  refining 
of  copper,  in  electroplating  and  electrotyping,  as  a  mordant  in 


434  A  TEXTBOOK  OF  CHEMISTRY 

dyeing  and  in  the  gravity  cells  formerly  much  used  for  tele- 
graphic purposes.  The  anhydrous  salt  is  sometimes  used  for 
drying  alcohol,  as  it  is  very  hygroscopic. 

Vitriols.  Vitriol  is  a  very  old  name  given  to  sulfates  because 
many  of  them  have  a  glassy  appearance.  It  is  rarely  used,  now, 
except  for  copper  sulfate,  CuSO4.5H2O,  or  blue  vitriol,  ferrous 
sulfate,  FeSO4.7  H2O,  or  green  vitriol,  zinc  sulfate,  ZnSO4.7H2O, 
or  white  vitriol,  and  for  sulfuric  acid,  oil  of  vitriol.  The  last 
name  is  derived  from  an  old  method  of  preparing  the  acid  by 
distilling  green  vitriol. 

*  Cupric  Nitrate,  Cu(NO3)2.6H2O,  is  formed  in  blue  tabular 
crystals  when  solutions  of  the  salt  are  crystallized  at  tempera- 
tures below  24.5°.  Above  that  temperature  the  hydrate 
Cu(NO3)2.3H2O  is  obtained.  On  heating  moderately  the  salt 
loses  water,  and  at  a  higher  temperature  it  is  decomposed  into 
copper  oxide,  CuO,  oxygen  and  nitrogen  dioxide,  NO2. 

Ammoniocupric  Sulfate,  CuSO4.H2O.4NH3.  When  ammonia 
is  added  to  a  solution  of  copper  sulfate,  a  very  intense  blue 
color  is  produced.  Alcohol  precipitates  from  such  a  solution 
ammoniocupric  sulfate.  This  may  be  considered  as  the  hydrate 
of  copper  sulfate  in  which  four  molecules  of  water  have  been 
replaced  by  four  molecules  of  ammonia..  Similar  compounds 
are  formed  with  other  salts  of  copper,  and  the  reaction  may  be 
used  for  the  detection  of  minute  quantities  of  the  element.  The 
formula,  [Cu(NH3)4]SO4.H2O,  used  by  Werner  and  others  brings 
out  more  clearly  the  intimate  relation  between  the  ammonia 
and  the  copper. 

*Cuprous  Cyanide,  Cu2(CN)2.  Solutions  of  copper  salts  react 
with  a  hot  solution  of  potassium  cyanide  in  very  much  the 
same  manner  as  with  potassium  iodide : 

2  CuS04  +  4  KCN  =  Cu2(CN)2  +  2  K2SO4  +  C2N2 

If  the  compounds  are  used  in  the  proportions  given  in  the 
equation,  the  cuprous  cyanide  separates  as  a  white  precipitate. 
If  an  excess  of  the  potassium  cyanide  is  used,  a  complex  salt, 
K3Cu(CN)4,  which  is  easily  soluble,  is  formed.  This  salt  gives 


ELECTROMOTIVE  SERIES  435 

a  solution  containing  so  few  copper  ions  that  no  precipitate  is 
formed  by  hydrogen  sulfide.  This  property  is  sometimes  used 
to  separate  copper  from  cadmium,  as  the  latter  is  precipitated 
as  cadmium  sulfide,  CdS,  under  similar  conditions. 

Precipitation  of  Copper  by  Iron,  Electromotive  Series.  If  an 
iron  nail  is  dipped  in  a  solution  of  copper  sulfate  or  copper 
chloride,  metallic  copper  will  be  deposited  and  iron  will  pass 
into  solution.  We  may  formulate  the  reaction  thus : 

Cu++  +  Fe  =  Fe++  +  Cu 

The  copper  ion,  Cu++,  loses  its  electrical  charge,  giving  it  to  the 
iron,  while  the  copper  assumes  the  metallic  form.  If  a  strip  of 
copper  and  one  of  iron  are  suspended  in  any  ordinary  electrolyte, 
an  electrical  current  will  pass  from  the  copper  to  the  iron  on 
connecting  the  two  metals  by  means  of  a  wire,  while  within  the 
electrolyte  the  positive  ions  will  travel  toward  the  copper  and 
the  negative  ions  toward  the  iron.  As  before,  the  iron  will 
pass  into  solution  and  copper  will  be  deposited  on  the  copper 
strip,  if  the  electrolyte  is  copper  sulfate.  On  the  basis  of  similar 
experiments  a  table  of  the  metals  may  be  arranged  in  such  a 
manner  that  each  metal  in  the  series  will  be  positive  toward  all 
of  the  metals  on  one  side  of  it  and  negative  toward  those  on  the 
other  side. 

The  simplest  method  of  looking  at  these  phenomena  is  to 
consider  that  each  metal  in  contact  with  a  solution  has  a  cer- 
tain solution-pressure  which  tends  to  cause  the  metal  to  pass 
into  solution.  We  may  suppose  that  a  few  atoms  do  actually 
leave  the  piece  of  metal  and  pass  into  solution  as  ions,  but  this 
would  give  the  solution  a  positive  charge  and  leave  the  metal 
negative ;  and  unless  some  means  is  provided  for  the  escape  of 
the  electrical  charges  the  electromotive  force  set  up  will  very 
quickly  balance  the  solution  pressure  and  the  process  will  cease. 
If  the  metal  is  dipped  in  a  solution  of  one  of  its  salts,  the  ions 
of  the  metal  in  solution  will  partly  or  wholly  counterbalance 
the  solution  pressure  of  the  metal,  the  quantitative  effect 
depending  on  the  concentration  of  the  ions  and  the  nature  of  the 


436 


A  TEXTBOOK  OF  CHEMISTRY 


metal.  For  many  metals  this  effect  may  exceed  the  solution 
pressure  of  the  metal  and  the  latter  will  assume  a  positive  poten- 
tial with  reference  to  the  solution  through  the  deposit  and  dis- 
charge of  some  metal  ions  on  the  plate. 

The  differences  in  potential  between  the  various  metals  and 
normal  solutions  of  their  ions  are  given  in  the  following  table : 1 

ABSOLUTE  POTENTIAL  OP  ELEMENTS  IN   CONTACT  WITH   NORMAL 
SOLUTIONS  OF  THEIR  SALTS.     ELECTROMOTIVE  SERIES 


ELEMENT 


ABSOLUTE 
POTENTIAL 


Electropositive  End 

Li -  2.740 

K -  2.644 

Na -  2.434 

Ba    (-2.6) 

Sr (-2.6) 

Ca (-2.4) 

Mg  .  .  .  .  -  1.31 

Al -1.04? 

Mn -  0.84 

Zn -  0.52 

S -  0.31 

Fe -  0.19 

Cd -  0.16 

Te    -  0.08 

Co..  -0.05 


ELEMENT 

Ni 

Pb   

Sn 

(H 

As 

Cu  (bivalent)    . 

Bi 

Cu  (univalent) , 

Sb 

Hg  (univalent) 

Pd    

Ag   

Pt 

Au   

F 

Cl 

Br 

I 

O 

Electronegative 


ABSOLUTE 
POTENTIAL 

+  0.02  ? 
+  0.12 
+  0.14 
+  0.24) 
+  0.53  ? 
+  0.58 
+  0.63  ? 
+  0.71 
+  0.71  ? 
+  0.99 
+  1.03? 
+  1.04 
+  1.10? 
+  1.7? 
(+2.1) 
(+  1-59) 
(+  1.32) 
(+  0.78) 
(+0.65) 
End 


The  values  in  parenthesis  have  not  been  measured  directly 
but  were  calculated  from  thermochemical  data.  Elements 
which  assume  a  high  negative  potential  in  contact  with  solu- 
tions of  their  salts  are  called  electropositive  because  the  ions 


Wilh.  Palmaer,  Nernst's  Festschrift,  p.  336  (1907). 


ELECTROMOTIVE   SERIES 


437 


which  separate  are  strongly  positive.     Elements  which  assume 
a  positive  potential  are  called  relatively  electronegative. 

This  table  may  be  used  in  two  ways :  first,  any  metal  of  the 
table  may  be  precipitated  by  any  other  metal  which  has  a 
lower  absolute  potential  and  it  will  precipitate  any  metal  with 
a  higher  potential  —  thus  copper  will  precipitate  silver,  but  it 
will  be  precipitated  by  lead  or  iron ;  second,  the  values  may  be 
used  to  calculate  the  electromotive  force  of  a  galvanic  jcell  in 
which  two  of  the  metals  are  used.  For  an  accurate  calculation 
it  is  necessary  to  take  account  of  the 
concentration  of  the  solution  in 
contact  with  each  electrode,  and 
when  two  solutions  are  used,  the 
difference  in  potential  between  the 
solutions  must  also  be  considered. 

The  common  gravity  cell  (Fig.  99) 
consists  of  a  copper  plate  A,  in  con- 
tact with  a  solution  of  copper  sul- 
fate,  and  a  zinc  plate  B,  in  contact 
with  a  solution  of  zinc  sulfate.  The 
electromotive  force  of  the  battery 
is  approximately  the  difference  be- 
tween the  absolute  potentials  of 
copper  and  zinc,  or  +  0.58—  (—0.52) 
=  1.10  volts. 

The  electromotive  force  of  the  Weston  cells,  consisting  of 
mercury  in  contact  with  mercurous  sulfate  and  cadmium  in 
contact  with  cadmium  sulfate,  would  be  for  normal  solutions  of 
each  +  0.99  -  (-0.16)  =  +  1.15. 

The  Clark  cell,  which  has  zinc  in  place  of  the  cadmium, 
would  have  for  normal  solutions  an  electromotive  force  of 
+  0.99  —  (—0.52)  .=  +  1.51.  In  both  cases,  however,  the 
electropositive  metal  (cadmium  or  zinc)  is  in  contact  with  a 
concentrated  or  saturated  solution  of  its  salt,  and  the  mercurous 
sulfate  is  only  very  slightly  soluble.  This  lessens  the  difference 
of  potential,  and  the  actual  value  for  the  Clark  cell  is  1.434 


Fig.  99 


438  A  TEXTBOOK  OF  CHEMISTRY 

volts,  and  for  the  Weston  cell,  in  which  a  saturated  solution  of 
cadmium  sulfate  is  used,  it  is  1.019  volts. 

Faraday's  Law.  If  the  same  electrical  current  is  passed 
through  a  series  of  cells  containing  electrolytes  which  are  de- 
composed by  the  current,  it  is  found  that  the  amounts  of  the 
elements  which  separate  at  the  electrodes  will  be  proportional 
to  the  equivalents  of  the  elements,  i.e.  to  the  atomic  weights  of 
the  elements  divided  by  their  valences.  In  a  series  of  cells  con- 
taining solutions  of  the  following  compounds  : 


AgNO3          CuCl  CuCl2          SnCl2  SnCl4 

H  O      Ag     O        Cu     Cl         Cu      Cl        Sn     Cl          Sn      Cl 
Ig8g    108g8g    63.6  g  35.5  g    31.8  g  35.5  g    59  g  35.5  g    29.5  g  35.5  g 

the  same  current  which  causes  the  liberation  of  one  gram  of 
hydrogen  will  liberate  8  grams  of  oxygen,  35.5  grams  of  chlorine, 
108  grams  of  silver,  63.6  grams  of  copper  from  cuprous  chloride, 
31.8  grams  of  copper  from  cupric  chloride,  59  grams  of  tin  from 
stannous  chloride  and  29.5  grams  of  tin  from  stannic  chloride. 
This  relation  was  discovered  by  Faraday  in  1834  and  is  known 
as  "Faraday's  Law."  The  rational  explanation  of  the  law  is 
that  each  univalent  ion  is  associated  with  a  unit  charge  of  elec- 
tricity, each  bivalent  ion  with  twice  this  unit  charge  and  a 
quadrivalent  ion  with  four  times  this  amount.  In  accordance 
with  the  electron  theory  this  unit  charge  is  the  charge  carried 
by  an  electron,  and  the  fact  that  it  does  not  seem  possible  to 
account  for  Faraday's  law  without  assuming  that  there  is  a 
definite,  unit  charge,  is  one  of  the  most  important  facts  sup- 
porting the  theory. 

It  should  not  be  overlooked  that  the  amount  of  energy  re- 
quired to  decompose  a  gram  equivalent  of  different  electrolytes 
is  not  the  same.  The  difference  in  potential  between  the  two 
electrodes  varies  from  cell  to  cell  ;  and  the  energy  consumed  in 
the  cell  depends  on  two  factors  —  the  quantity  of  electricity 
passing  through  the  cell  and  the  fall  in  potential  from  one  elec- 
trode to  the  other.  The  latter  depends  on  the  absolute  poten- 
tials of  the  substances  separating  at  the  two  electrodes  and  on 


SILVER  439 

the  concentration  of  the  electrolyte  exactly  as  in  the  galvanic 
cells  discussed  in  the  preceding  paragraph.  An  electrolytic  cell 
may  be  considered  as  a  galvanic  battery  in  which  the  direction 
of  the  current  has  been  reversed  by  the  application  of  an  external 
electromotive  force. 

Silver,  Ag,  107.88.  Copper,  silver,  gold  and  all  metals  which 
are  more  electronegative  than  hydrogen  in  the  electromotive 
series  are  found  sometimes  in  the  free  state  in  nature.  The 
only  other  metals  found  in  the  free  state  are  iron,  cobalt  and 
nickel,  and  possibly  some  others  which  have  come  to  the  earth 
in  the  form  of  meteorites  and  have  not  had  time  since  their 
arrival  to  become  completely  oxidized. 

Silver  is  also  found  as  the  sulfide,  Ag2S,  either  alone,  or  more 
often  associated  with  other  sulfides  and  especially  with  lead 
sulfide  or  galena,  PbS.  It  is  also  found  as  the  chloride  in  the 
mineral  cerargyrite,  or  horn  silver,  AgCl. 

Metallurgy.  In  the  electrolytic  refining  of  copper  a  chloride 
is  added  to  the  electrolyte,  and  this  causes  the  silver  to  separate 
as  the  insoluble  chloride,  Ag;Cl,  with  the  slimes  from  which  the 
silver  and  gold  are  recovered. 

The  metallic  lead  obtained  from  galena,  PbS,  always  contains 
some  silver.  This  is  recovered  either  by  Pattison's  process  of 
crystallization  or  by  Parke's  process  of  extraction  with  zinc. 

*  Pattinson's  Process  depends  on  the  principle  that  a  solution 
melts  at  a  lower  temperature  than  the  solvent.  If  lead  contain- 
ing'silver  is  melted  in  an  iron  pot,  on  cooling,  crystals  of  nearly 
pure  lead  separate  at  first,  leaving  a  solution  or  alloy  of  silver 
and  lead  which  is  richer  in  silver  than  the  original  metal.  The 
crystals  of  lead  are  skimmed  out  and  transferred  to  another  pot 
on  one  side  while  the  richer  alloy  is  transferred  to  a  pot  on  the 
opposite  side.  By  repeating  the  operation  several  times,  nearly 
pure  lead  is  obtained  at  one  end  of  the  series  of  pots  and  an 
alloy  comparatively  rich  in  silver  at  the  other  end.  This  rich 
alloy  is  then  heated  in  a  furnace  with  free  access  of  air  till  the 
lead  is  oxidized  to  litharge,  PbO,  leaving  very  nearly  pure 
silver  behind.  While  this  process  is  no  longer  used,  it  is  of 


440 


A  TEXTBOOK  OF  CHEMISTRY 


historical  interest  and  also  interesting  because  of  the  principles 
of  crystallization  involved. 

Cupellation,  Assaying.  The  process  of  oxidizing  lead  contain- 
ing silver  is  often  carried  out  on  a  small  scale  in  a  muffle  furnace 
(Fig.  100)  on  a  small  cup  of  porous  bone  ash,  called  a  cupel, 

which  absorbs  the 
litharge.  Small 
amounts  of  other 
metals  are  oxi- 
dized with  the 
lead  and  absorbed 
by  the  cupel  so 
that  an  almost 
pure  silver  bead 
remains.  This 
can  be  weighed 
as  a  means  of 
determining  the 
amount  of  silver 
in  lead  bullion. 
The  gold  and  sil- 
ver of  an  ore  may 
be  concentrated 
in  a  lead  button 
which  can  then  be 
cupelled.  Such 
determinations, 


MUFFLE  FURNACE 

Fig.  100 


or     assays, 


fur- 


nish the  basis  for 
commercial  trans- 
actions with  ores 
and  bullion. 

Parke's  Proc- 
ess for  the  ex- 
traction of  silver  from  lead  is  exactly  analogous  to  the  extraction 
of  a  substance  from  water  by  means  of  ether.  When  zinc  and 


SILVER  441 

lead  are  melted  together,  on  allowing  the  mixture  to  stand  for 
a  few  minutes  in  the  melted  condition,  nearly  all  of  the  zinc 
will  rise  to  the  top  in  the  form  of  a  solution  of  lead  in  zinc, 
containing  only  a  small  quantity  of  lead.  The  solution  of 
zinc  in  lead  below  will  contain  only  a  small  amount  of  zinc. 
The  silver  and  gold,  however,  are  much  more  soluble  in  zinc 
than  in  lead,  and  a  comparatively  small  amount  of  zinc  will 
carry  nearly  all  of  the  silver  to  the  top  with  it.  If  much  silver 
is  present,  it  may  be  desirable  to  repeat  the  process.  The 
alloy  of  zinc  and  silver,  which  is  skimmed  from  the  top,  may 
be  placed  in  a  retort  and  the  zinc  distilled  away.  The  residual 
lead  is  then  removed  by  cupellation. 

*  Amalgamation  Process.  From  ores  containing  little  lead  or 
copper,  silver  was  formerly  often  separated  by  an  amalgamation 
process.  The  ore  was  pulverized  and  intimately  mixed  with 
metallic  mercury  and  water.  The  mercury  dissolved,  or  amal- 
gamated with  the  silver,  if  that  was  in  the  free  state.  If  present 
as  the  chloride,  the  mercury  reduced  it  to  the  metallic  state : 

AgCl  +  Hg  =  Ag  +  HgCl 

Mercurous 
Chloride 

If  the  silver  was  present  as  the  sulfide,  it  might  be  reduced  by 
adding  iron  turnings,  or  the  ore  was  roasted  with  salt  before  it 
was  amalgamated : 

2  NaCl  +  Ag2S  =  Na2S  +  2  AgCl 

After  the  amalgamation  was  complete  the  lighter  materials  of 
the  ore  were  washed  away  from  the  amalgam  in  a  current  of 
water,  the  amalgam  was  collected,  and  the  larger  portion  of  the 
mercury  removed  by  squeezing  it  through  chamois  skin,  which 
allows  mercury  but  not  the  solid  amalgam  to  pass.  The  latter 
was  heated  in  a  retort  to  remove  the  rest  of  the  mercury. 

Other  Processes  for  the  Recovery  of  Silver.  The  recovery  of 
silver  in  the  electrolytic  refining  of  copper  has  been  referred  to 
above.  The  cyanide  process  (p.  446)  has  largely  displaced  all 
others  for  obtaining  both  silver  and  gold. 


442  A  TEXTBOOK  OF  CHEMISTRY 

The  production  of  silver  in  the  United  States  in  1910  was 
57,000,000  troy  ounces,  worth  $31,000,000. 

Properties  of  Silver.  Alloys.  Silver  is  a  soft,  white  metal, 
extremely  malleable  and  ductile.  It  is  the  best  conductor  of 
electricity  known.  (See  table  at  the  end  of  this  chapter.)  It 
is  also  a  very  good  conductor  of  heat,  the  two  properties  being 
closely  parallel  in  most  cases,  a  fact  which  is  readily  explained 
by  the  electron  theory. 

Silver  is  too  soft  for  satisfactory  use  in  the  free  state  and  is 
usually  alloyed  with  copper  to  harden  it  for  the  manufacture  of 
coins  or  of  table  ware.  "  Sterling  "  silver,  used  in  British  coins, 
contains  1\  per  cent  of  copper.  The  coins  of  the  United  States 
contain  10  per  cent  of  copper.  Silver  does  not  tarnish  in  moist 
or  dry  air,  but  it  is  easily  blackened  by  hydrogen  sulfide  or  by  a 
solution  of  a  sulfide. 

Silver  dissolves  easily  in  nitric  acid  or  in  hot  concentrated 
sulfuric  acid,  with  evolution  of  nitric  oxide,  NO,  in  the  former 
case  and  of  sulfur  dioxide  in  the  latter.  It  is  not  attacked  by 
hydrochloric  acid,  partly  because  of  the  insolubility  of  the 
chloride. 

Silver  Plating.  When  silver  is  deposited  on  a  cathode  from 
a  solution  of  silver  nitrate,  or  of  some  other  silver  salt  which  is 
ionized  largely  in  solution,  it  assumes  a  crystalline  form  and 
gives  a  surface  which  is  not  suitable  for  plated  ware.  If  a  solu- 
tion of  potassium  silver  cyanide,  KAg(CN)2,  is  used  in  place  of 
one  of  the  ordinary  salts,  a  smooth,  coherent  deposit  can  be 
obtained.  In  the  electrolysis  the  object  to  be  plated  is  made  the 
cathode  while  a  plate  of  pure  silver  is  used  as  an  anode.  The 
object  to  be  plated  requires  very  careful  cleaning,  as  the  slightest 
film  of  grease  will  cause  the  deposited  silver  to  flake  off. 

Silver  Oxide,  Ag2O,  is  formed  when  a  solution  of  sodium 
hydroxide  is  added  to  a  solution  of  silver  nitrate.  It  may  be 
supposed  that  silver  hydroxide,  AgOH,  is  formed  at  first,  but 
that  this  immediately  decomposes  into  silver  oxide  and  water. 
The  decomposition  seems  to  represent  an  equilibrium,  however, 
which  lies  far  on  the  side  toward  the  formation  of  the  oxide, 


SILVER  443 

since  the  oxide  dissolves  slightly  in  water  and  the  solution  has 
an  alkaline  reaction.  Also,  silver  oxide  cannot  be  completely 
freed  from  water  below  the  temperature  at  which  it  dissociates 
rapidly  into  silver  and  oxygen.  The  solution  of  silver  hydroxide 
is  ionized  to  the  extent  of  70  per  cent,  indicating  that  the  com- 
pound is  a  comparatively  strong  base. 

Silver  oxide  is  easily  decomposed  by  heat,  but  this  reaction, 
also,  is  reversible  :  2  ^0^4  Ag  +  Q2 

As  in  all  cases  of  an  equilibrium  between  a  gaseous  and  a 
solid  phase,  the  equilibrium  depends  only  on  the  temperature 
and  the  pressure  of  the  gaseous  component  and  not  upon  the 
relative  amounts  of  the  two  phases,  because  the  reaction  can 
occur  only  at  the  surface  between  the  gas  and  the  .solid,  and  not 
throughout  the  mass  of  either  phase.  At  302°  the  dissociation 
pressure  of  silver  oxide  is  20.5  atmospheres.  If  the  pressure  is 
decreased,  more  of  the  oxide  will  decompose  until  all  is  decom- 
posed or  the  pressure  rises  to  20.5  atmospheres  again.  If  the 
pressure  is  increased,  oxygen  will  combine  with  the  silver  till 
all  is  converted  into  the  oxide  or  till  the  pressure  falls  to  20.5 
atmospheres.  At  325°  the  dissociation  pressure  is  32  atmos- 
pheres ;  at  445°  it  is  207  atmospheres.  The  dissociation  pres- 
sure of  silver  oxide  is  0.2  of  an  atmosphere  at  121°.  As  0.2  of 
an  atmosphere  is  the  partial  pressure  of  oxygen  in  the  air,  it 
follows  that  at  temperatures  above  121°,  silver  oxide  will  de- 
compose completely  if  exposed  to  the  air,  while  at  temperatures 
below  that  metallic  silver  would  change  to  silver  oxide.  At 
that  temperature,  however,  both  the  oxidation  and  its  decom- 
position occur  very  slowly  indeed,  if  no  catalyzer  is  present 
(Lewis,  J.  Am.  Ch.  Soc.  28,  139  (1906)). 

Molten  silver  absorbs  oxygen,  which  it  gives  out  again,  in 
part,  as  the  metal  solidifies.  It  seems  probable  that  the  absorbed 
oxygen  is  not  in  chemical  combination  with  the  silver. 

*  Silver  Peroxide,  Ag2O2,  is  formed  as  a  brown  or  black  coat- 
ing by  the  action  of  ozone  on  silver.  It  is  also  deposited  on  the 
anode  as  a  black,  crystalline  compound  in  the  electrolysis  of 


444  A  TEXTBOOK  OF  CHEMISTRY 

silver  nitrate  with  an  electrical  pressure  of  about  15  volts  and 
the  use  of  a  diaphragm  of  porous  porcelain. 

Silver  Nitrate,  AgNOs.  This  is  easily  prepared  by  dissolving 
silver  in  nitric  acid  and  evaporating  the  solution  till  the  salt 
will  crystallize  on  cooling.  It  forms  tabular  crystals  of  the 
rhombic  system.  It  is  easily  soluble  in  water  and  is  used  as 
the  starting  point  for  the  preparation  of  most  of  the  other  com- 
pounds of  silver.  It  melts  at  208.6°  and  is  sometimes  cast  in 
small  sticks  for  use  as  a  cauterizing  and  germicidal  agent  in 
medicine.  In  connection  with  this  use  an  old  name,  lunar 
caustic,  is  still  employed  —  a  name  which  comes  to  us  from  the 
alchemists,  who  recognized  a  symbolical  relation  between  silver 
and  the  moon  (Latin,  luna). 

*  Silver  Nitrite,  AgNO2,  is  a  white,  difficultly  soluble  salt 
easily  prepared  by  precipitating  a  solution  of  silver  nitrate 
with   a  moderately  concentrated  solution  of  sodium  nitrite.     It 
is  used  in  water  analyses  for  the  preparation  of  standard  solu- 
tions of  nitrites. 

*  Silver  Sulfate,  Ag2SO4,  is  prepared  by  warming  silver  with 
concentrated  sulfuric  acid.     The  salt  is  only  moderately  soluble 
(1 :  200  in  cold  water).     It  is  sometimes  used  to  remove  chlorine, 
bromine  or  iodine  from  solutions  when  the  presence  of  a  nitrate 
is  to  be  avoided. 

Silver  Chloride,  AgCl,  Silver  Bromide,  AgBr,  and  Silver 
Iodide,  Agl,  are  almost  insoluble  compounds  which  separate  as 
curdy  precipitates  on  adding  one  of  the  halogen  acids  or  a  halide 
to  a  solution  of  silver  nitrate.  The  chloride  is  white ;  the  bromide, 
yellowish  white ;  and  the  iodide,  yellow.  All  three  of  the  salts 
are  sensitive  to  light  and  lose  chlorine,  bromine  or  iodine, 
slowly  in  diffused  light,  rapidly  in  sunlight  or  in  artificial  light 
which  contains  rays  of  short  wave  lengths  —  those  of  the  violet 
end  of  the  spectrum  or  beyond.  The  action  seems  to  be  closely 
related  to  the  effect  of  sunlight  or  the  magnesium  light  in  causing 
the  combination  of  hydrogen  and  chlorine  (p.  105). 

Photography.  Photographic  "  dry  "  plates  are  prepared  by 
spreading  a  thin  film  of  an  emulsion  of  silver  bromide  in  a  solu- 


PHOTOGRAPHY.    GOLD  445 

tion  of  gelatin  over  a  plate  of  glass  or  of  transparent  celluloid 
and  allowing  the  film  to  dry  in  a  dark  room.  When  such  a  film 
has  been  exposed  to  the  image  formed  by  the  lens  of  a  camera 
no  change  in  the  appearance  of  the  film  can  be  noticed,  but  if 
the  exposed  plate  is  placed  in  a  solution  of  hydroquinone,  pyro- 
gallol  or  some  other  reducing  agent  used  as  a  "  developer/' 
those  portions  of  the  silver  bromide  which  were  struck  by  the 
light  will  be  reduced  to  metallic  silver,  which  appears  dark  and 
opaque,  while  the  portions  under  the  dark  parts  of  the  image  of 
the  camera  will  not  be  affected.  When  the  lights  and  shadows 
have  been  brought  out  sufficiently  by  the  developer,  the  plate  is 
placed  in  a  solution  of  sodium  thiosulfate,  Na2S2Os,  which  dis- 
solves silver  bromide  readily,  while  the  metallic  silver  is  not 
affected.  If  the  unchanged  silver  bromide  were  not  removed, 
it  would  darken  later  on  exposure  to  light,  and  the  picture  would 
be  destroyed.  The  process  of  removing  the  unchanged  silver 
bromide  is  called  "  fixing  "  the  picture.  The  picture  obtained 
is  dark  where  the  object  photographed  was  light  and  light 
where  the  object  was  dark.  For  this  reason  it  is  called  a  "  nega- 
tive." 

In  order  to  obtain  a  "  positive,"  which  will  reproduce  the 
lights  and  shadows  of  the  object,  the  negative  is  placed  over  a 
paper  which  has  been  coated  with  a  film  of  silver  chloride  or 
bromide  and  is  exposed  to  the  light  for  "  printing."  The  light 
passing  through  the  light  portions  of  -the  negative  darkens  the 
silver  salt  beneath,  while  other  portions  of  the  salt  are  protected 
by  the  opaque  portions  of  the  negative.  The  picture  on  the 
paper  must  be  fixed  by  means  of  sodium  thiosulfate  as  is  done 
with  the  negative.  It  may  also  be  "  toned  "  by  immersion  in  a 
solution  of  gold  trichloride,  AuCls,  or  of  chloroplatinic  acid, 
H2PtCle,  which  will  cause  the  silver  to  be  replaced  by  gold  or 
platinum.  The  gold  gives  a  reddish  tone,  the  platinum,  a  steel- 
gray  color. 

Gold,  Au.,  197.2,  is  found  almost  always  in  the  free  state  in 
nature,  very  rarely  in  large  nuggets  weighing  several  pounds, 
usually  in  small  grains  mixed  with  sand  or  gravel  or  dissemi- 


446  A  TEXTBOOK  OF  CHEMISTRY 

nated  through  quartzite,  granite,  pyrite  and  other  rocks  and 
minerals.  In  Colorado  and  in  some  other  localities  it  is  some- 
times found  combined  with  or  alloyed  with  tellurium  in  a 
mineral  having  approximately  the  composition  AuTe2.  Al- 
though gold  is  almost  always  found  mixed  with  very  large 
quantities  of  worthless  or  comparatively  worthless  minerals,  it 
is  very  widely  diffused  and  a  careful  examination  reveals  traces 
of  gold  in  almost  any  rock,  soil,  clay  or  other  natural,  inorganic 
mixture  which  is  tested.  It  is  claimed  that  sea  water  contains 
from  0.03  to  0.06  gram  of  gold  in  a  ton  (Liversidge,  see  Chem. 
Centralblatt,  1905,  II,  648).  This  would  be  worth  from  two 
to  four  cents,  but  the  total  amount  of  gold  in  the  ocean  is,  of 
course,  very  large. 

Metallurgy.  Gold  has  a  specific  gravity  of  19.26,  while 
that  of  ordinary  minerals  averages  about  2.6.  Such  minerals 
may,  therefore,  be  separated  from  the  gold  by  "  washing  "  in  a 
current  of  water,  which  carries  the  lighter  minerals  away,  leaving 
the  gold.  For  rich  sands  or  gravels  the  process  may  sometimes 
be  carried  out  by  hand  in  a  pan  —  an  operation  which  has  given 
the  well-known  expression  "  pan  out."  On  a  large  scale,  in 
"  hydraulic  mining,"  masses  of  sand  and  gravel  are  washed 
away  by  powerful  streams  of  water,  the  material  running 
through  sluiceways  in  which  are  placed  crosspieces  to  retain 
the  gold.  Metallic  mercury  is  usually  placed  back  of  the  cross- 
pieces,  or  riffles.  This  amalgamates  with  the  gold  and  retains 
it.  Dredging  is  also  used  to  get  gold-bearing  material  into 
sluices. 

Cyanide  Process.  A  few  years  ago  large  quantities  of  gold 
were  obtained  by  a  "  chlorination  process,"  in  which  the  gold 
was  dissolved  by  chlorine  obtained  from  bleaching  powder  and 
sulfuric  acid.  This  process  has  been  almost  completely  replaced 
by  the  cyanide  process.  The  cyanide  process  has  also  displaced 
smelting  processes  in  some  cases. 

In  the  presence  of  air  to  furnish  oxygen,  gold  dissolves  in  a 
solution  of  potassium  cyanide  as  potassium  aurous  cyanide, 
KAu(CN)2.  Metallic  silver  and  silver  chloride  will  also  dissolve. 


GOLD  447 

4  Au  +  8  KCN  +  O2  +  2  H2O  =  4  KAu(CN)2  +  4  KOH 
4  Ag  +  8  KCN  +  02  +  2  H2O  =  4  KAg(CN)2  +  4  KOH 

The  necessity  of  air  in  the  solution  can  be  shown  by  shaking 
gold  leaf  with  a  solution  of  potassium  cyanide  which  is  free  from 
oxygen  or  through  which  a  current  of  hydrogen  is  passed.  It 
will  not  dissolve,  while  it  dissolves  readily  on  passing  a  current 
of  air.  Potassium  permanganate,  potassium  chromate,  sodium 
peroxide,  nitrobenzene  or  some  other  oxidizing  agent  is  sometimes 
used  to  assist  in  the  solution  of  the  gold.  The  gold  is  precipi- 
tated from  the  solution  by  means  of  zinc : 

2  KAu(CN)2  +  Zn  =  K2Zn(CN)4  +  2  Au 

The  solution  containing  potassium  zinc  cyanide  may  be  used 
for  a  new  lot  of  ore. 

Native  gold  frequently  contains  silver,  and  gold  is  often 
obtained  along  with  silver  from  silver  or  copper  ores.  From 
such  alloys  the  silver  can  be  removed  by  solution  in  dilute  nitric 
acid  or  in  concentrated  sulfuric  acid,  provided  not  more  than 
35  per  cent  of  gold  is  present.  Alloys  which  are  richer  in  gold 
than  this  are  melted  with  enough  silver  to  reduce  the  amount  of 
gold  to  one  third  or  one  fourth  of  the  whole.  The  separation 
with  concentrated  sulfuric  acid  can  be  carried  out  in  cast-iron 
kettles  and  is  cheaper  than  that  with  nitric  acid.  It  also  has 
the  advantage  that  any  platinum  which  is  present  remains  with 
the  gold,  while  small  amounts  of  platinum  dissolve  with  the  silver 
in  nitric  acid.  The  separation  of  gold  from  silver  by  solution  of 
the  silver  is  called,  technically,  "  parting." 

The  production  of  gold  in  the  United  States  in  1910  was 
4,657,018  troy  ounces  worth  $96,269,100.  The  coining  value 
of  an  ounce  of  gold  is  $20.67183.  The  average  yearly  production 
of  gold  in  the  world  during  the  first  half  of  the  nineteenth  cen- 
tury was  only  about  800,000  troy  ounces.  The  present  annual 
production  in  the  United  States  alone  is  nearly  six  times  that 
and  the  annual  production  in  the  world  is  many  times  that  of 
70  or  80  years  ago. 


448  A  TEXTBOOK  OF  CHEMISTRY 

Properties  of  Gold.  Gold  is  yellow  by  reflected  light  but 
transmits  the  complementary  color,  green,  through  very  thin 
films,  as  through  gold  leaf  held  between  two  glass  plates.  It  is 
the  most  ductile  and  malleable  of  all  the  metals  and  can  be 
drawn  into  exceedingly  fine  wires  and  beaten  into  very  thin 
leaves.  Its  specific  gravity  is  about  19.3,  varying  considerably 
with  the  method  of  preparation  and  treatment.  It  melts  at 
1063°.  "  Its  electrical  conductivity  is  two  thirds  that  of  silver. 

Gold  is  insoluble  in  any  one  of  the  common  acids,  alone,  but 
it  dissolves  readily  in  aqua  regia,  a  mixture  of  three  volumes  of 
hydrochloric  acid  with  one  of  nitric.  Its  solubility  in  chlorine 
water  has  been  mentioned  above.  It  dissolves  also  in  selenic 
acid. 

Alloys  of  Gold.  Gold  is  a  very  soft  metal  and  is  alloyed  with 
other  metals,  especially  with  silver  and  copper  for  use  in  jewelry 
and  coins.  British  gold  coins  are  22  carats  fine,  i.e.  ff  pure 
gold.  In  reference  to  gold  the  term  "  carat  "  is  used  to  designate 
the  number  of  parts  in  a  total  of  24  which  consist  of  pure  gold. 
Thus  18-carat  gold  is  yf  or  J  pure  gold.  American  gold 
coins  are  made  on  a  decimal  basis  "  900-fine,"  i.e.  900  parts  in 
1000  are  pure  gold.  The  American  eagle  contains  900  parts  of 
gold  and  100  parts  of  copper. 

*  Oxides  of  Gold.     Three  oxides  of  gold  have  been  described : 
gold  monoxide,  Au2O;   gold  dioxide,  AuO;   and  gold  trioxide, 
Au2Oa.     There  seems  to  be  some  doubt  whether  the  second  of 
these  has  been  prepared  as  a  definite  compound.     Each  oxide  is 
decomposed  at  a  comparatively  low  temperature  into  gold  and 
oxygen. 

M 

*  Gold  Hydroxide,  Au— O — H,  may  be  prepared  by  adding 
magnesium  carbonate,  MgCOs,  to  a  solution  of  chloroauric  acid 
and  dissolving  the  excess  with  dilute  nitric  acid : 

HAuCl4  +  2  MgCO3  =  HAuO2  +  2  MgCl2  +  2  CO2 

Gold  hydroxide  is  a  yellowish  brown  precipitate  having  the 
properties  of  both  a  base  and  an  acid.  It  dissolves  in  hydro- 


ELECTRICAL  CONDUCTANCES 


449 


SPECIFIC  CONDUCTANCE  AND  RESISTANCE  OF  COMMON  METALS* 


TEMPER- 
ATURE 

SPECIFIC 
CONDUCTANCE 

SPECIFIC 
RESISTANCE 

RESISTANCE  IN 
OHMS  OF  WIRE 
1  M.  LONG  1  MM. 
IN  DIAMETER 

Aluminium    .     . 

0° 

35.6      X  104 

2.81  X  10~6 

0.036 

Bismuth    .     .     . 

18° 

0.82    X  104 

125.  X  10~6 

1.58 

Cadmium       .     . 

0° 

14.6      X  104 

6.85  X  10"6 

0.087 

Cobalt       .     .     . 

20° 

10.3      X  104 

9.7    X  10~6 

0.123 

Copper      .     .     . 

25° 

58.6      X  104 

1.71  X  HT6 

0.022 

Gold     .... 

0° 

47.5      X  104 

2.10  X  10~6 

0.027 

Iron  (electrolytic) 

— 

8.27    X  104 

12.1    X  10~6 

0.154 

Iron  (steel,  1%C) 

18° 

5.02    X  104 

19.9    X  10'6 

0.254 

Lead     .... 

0° 

5.14    X  104 

19.4    X  10~6 

0.247 

Magnesium    .     . 

0° 

23.0      X  104 

4.35  X  10'6 

0.055 

Mercury    .     .     . 

0° 

1.063  X  104 

94.07  X  10~6 

1.197 

Molybdenum 

25° 

17.9      X  104 

5.6    X  HT6 

0.071 

Nickel       .     .     . 

— 

11.1      X  104 

9.1    X  HT6 

0.116 

Palladium      .     . 

0° 

9.47    X  104 

10.6    X  10~6 

0.135 

Platinum  .     .     . 

20° 

10.7      X  104 

9.3    X  10-6 

0.118 

Potassium      .     . 

18° 

14.9      X  104 

6.7    X  10-6 

0.086 

Silicon       .     .     . 

— 

1.68    X  104 

59.5    X  10~6 

0.76 

Silver    .... 

25° 

61.74    X  104 

1.62  X  HT6 

0.020 

Sodium      .     .     . 

18° 

20.8      X  104 

4.8    X  10~6 

0.061 

Tin       .... 

18° 

8.3      X  104 

12.3    X  10~6 

0.169 

Tungsten      (fila- 

ment)    .     .     . 

20° 

17.9      X  104 

5.6    X  1(T6 

0.071 

Zinc      .... 

16° 

16.6      X  104 

6.0    X  lO'6 

0.076 

1  Specific  conductance  is  the  current  in  amperes  which  would 
pass  between  opposite  faces  of  a  cubic  centimeter  of  the  metal  under 
a  potential  difference  of  one  volt.  The  specific  resistance  is  the 
reciprocal  of  the  specific  conductance.  If  the  factor  10~6  is  omitted, 
the  specific  resistance  is  given  in  microhms,  i.e.  in  millionths  of 
an  ohm.  The  resistance  of  a  wire  1  m.  long  and  1  mm.  in  diameter 

is  found  by  multiplying  the  specific  resistance  by  10000  X 


3.1416 


450  A  TEXTBOOK  OF  CHEMISTRY 

chloric  acid  with  regeneration  of  chloroauric  acid.  It  also  dis- 
solves in  a  solution  of  potassium  hydroxide  and  from  the  solution 
potassium  aurate,  KAuO2.3  H2O,  may  be  crystallized. 

Chlorides  of  Gold.  Corresponding  to  the  three  oxides  there 
are  three  chlorides  of  gold :  gold  monochloride,  AuCl ;  gold 
dichloride,  AuCl2 ;  and  gold  trichloride,  AuCls.  The  last  of  these 
combines  with  hydrochloric  acid  to  form  chloroauric  acid, 
HAuCU,  and  this  last  compound  is  formed  when  gold  is  dissolved 
in  aqua  regia.  It  is  deposited  by  evaporating  and  cooling  such 
a  solution,  in  the  form  of  yellow  needles  having  the  composition 
HAuCl4.4  H2O.  It  is  a  monobasic  acid,  which  forms  well- 
defined,  crystalline  salts  with  many  of  the  metals  and  especially, 
also,  with  many  organic  bases.  The  potassium  salt, 
KAuCl4.2H2O,  the  calcium  salt,  Ca(AuCl4)2.6  H2O,  the 
ammonium  salt,  2  NH4AuCl4.5  H2O,  and  the  strychnine  salt, 
C2iH22N2O2.HAuCl4,  are  typical  of  such  compounds. 

EXERCISES 

1.  Write  the  equation  representing  the  action  of  nitric  acid  on  copper ; 
also  for  the  action  of  nitric  acid  on  cupric  sulfide.     Free  sulfur  and  nitric 
oxide  are  formed  in  the  latter  case. 

2.  Write  the  equations  for  the  action  of  nitric  acid  and  of  sulfuric 
acid  on  silver. 

3.  How  many  cubic  centimeters  of  nitric  acid  of  specific  gravity  1.42, 
containing  70  per  cent  of  the  pure  acid,  will  be  required  to  dissolve  a 
pound  (453  grams)  of  silver  ? 

4.  How  many  grams  of  concentrated  sulfuric  acid  of  98  per  cent  will 
be  required  to  dissolve  a  pound  of  silver,  taking  account  only  of  that 
which  enters  into  the  reaction  ? 

5.  Write  the  equation  for  the  action  of  potassium  iodide  on  copper 
sulfate;   also  for  the  reaction  between  sodium  thiosulfate  and  iodine. 
How  many  milligrams  of  crystallized  sodium  thiosulfate,  Na2S2O3.5  H2O, 
are  equivalent  to  63.6  milligrams  of  copper  in  these  reactions  ?    How 
many  milligrams  are  equivalent  to  one  milligram  of  copper  ? 

6.  Write  the  equation  representing  the  solution  of  gold  in  selenic 
acid,  H2SeO4.     The  gold  dissolves  in  the  trivalent  form,  while  some  of 
the  acid  is  reduced  to  selenious  acid. 


CHAPTER  XXVI 

GROUP  II,  ALKALI-EARTH  METALS:  BERYLLIUM,   CALCIUM, 
STRONTIUM,  BARIUM,  RADIUM 

THE  elements  calcium,  strontium  and  barium  are  called  alkali- 
earth  metals,  a  somewhat  indefinite  designation  coming  down  to 
us  from  the  time  of  the  alchemists  and  referring  to  the  fact  that 
the  hydroxides  of  these  metals  are  strong  bases.  The  elements 
beryllium  and  radium  are  not  included  under  the  designation. 
All  of  the  elements  of  Group  II  are  bivalent.  Calcium,  stron- 
tium, and  barium  decompose  water  at  ordinary  temperature, 
though  much  less  rapidly  than  the  alkali  metals.  Calcium 
hydroxide  is  difficultly  soluble  in  water.  Strontium  and  barium 
hydroxides  are  somewhat  more  soluble.  All  three  carbonates 
are  nearly  insoluble  in  water,  but  dissolve  as  bicarbonates  in 
water  containing  carbon  dioxide.  Barium  sulfate  is  one  of  the 
most  insoluble  salts,  while  calcium  sulfate  is  slightly  soluble 
(1  part  in  500  of  water). 

*  Beryllium,1  Be,  9.1.  The  mineral  beryl  is  a  silicate  of  beryl- 
lium and  aluminium,  having  the  composition  BesA^SieOig  (or 
3  BeO.Al 203.6  SiCy.  As  a  precious  stone,  different  forms  of 
the  mineral  are  known  as  emerald  and  aquamarine.  The  free 
element  is  silver-white  and  has  a  specific  gravity  of  1.9.  Both 
as  an  element  and  in  its  compounds  beryllium  resembles  mag- 
nesium and  aluminium  rather  than  calcium.  The  hydroxide, 
Be  (OH)  2,  is  nearly  insoluble  in  water  but  dissolves  both  in  acids 
and  in  alkalies,  exhibiting  in  this  way  both  acid  and  basic 
properties.  The  chloride,  BeCl2.4  H2O,  sulfate,  BeSO4.4  H2O, 

1  Sometimes  called  glucinum,  Gl,  a  name  which  has  some  claim 
to  priority.  The  name  beryllium  has  the  advantage  of  referring  to 
the  most  important  mineral  containing  the  element. 

451 


452  A  TEXTBOOK  OF  CHEMISTRY 

and  nitrate,  Be(NO3)2.3  H2O,  are  soluble  in  water.  The  car- 
bonate is  unstable  and  readily  loses  carbon  dioxide,  forming  a 
basic  carbonate. 

Calcium,  Ca,  40.07.  Occurrence.  Calcium  is  one  of  the  most 
abundant  and  most  important  of  the  elements.  It  is  found  as 
the  carbonate,  CaCO3,  in  the  minerals  calcite  and  aragonite,  in 
marble  and,  less  pure,  in  limestones  ;  as  the  sulfate  in  the  mineral 
gypsum,  CaSO4.2H2O;  as  the  phosphate,  Ca3(PO4)2,  in  bone 
ash  and  mineral  phosphates  and  in  the  mineral  apatite, 
Ca5(PO4)3F ;  and  as  the  fluoride  in  the  mineral  fluorite,  CaF2. 

Preparation,  Properties.  Metallic  calcium  can  be  prepared 
most  easily  by  the  electrolysis  of  the  fused  chloride  (E.  F.  Smith 
and  Goodwin,  J.  Am.  Chem.  Soc.  25,  873  (1903)).  It  is  a  white, 
crystalline  metal  which  decomposes  water  readily  at  ordinary 
temperatures,  with  the  formation  of  the  hydroxide,  Ca(OH)2. 
It  burns  in  air,  forming  both  the  oxide,  CaO,  and  the  nitride, 
Ca3N2.  The  latter  is  hydrolyzed  by  water  with  the  formation  of 
ammonia : 

Ca3N2  +  6  HOH  =  3  Ca(OH)2  +  2  NH3 

Calcium  melts  at  800°. 

*  Calcium  Hydride,  CaH2,  is  a  white,  crystalline  compound 
which  can  be  prepared  by  the  direct  union  of  the  elements. 

Calcium  Oxide,  CaO,  or  Lime  is  manufactured  on  a  large 
scale  by  heating  the  carbonate  in  lime  kilns.  In  the  older  forms 
a  mixture  of  limestone  and  fuel  was  placed  in  the  kiln.  The 
fuel  was  then  set  on  fire  and  allowed  to  burn  till  it  was  all  con- 
sumed. In  the  newer  forms  a  mixture  of  limestone  and  coal  is 
charged  into  the  kiln  at  the  top,  while  the  "  burnt "  lime  is 
removed  at  the  bottom,  without  stopping  the  process.  The 
dissociation  of  the  carbonate : 

CaCO3  ^±  CaO  +  CO2. 

is  a  reversible  reaction.  As  there  is  a  gaseous  constituent,  the 
system  has  a  characteristic  pressure  for  each  temperature. 
This  may  be  determined  by  heating  calcium  carbonate  in  a  plat- 
inum bulb  connected  with  a  manometer.  The  pressure  of  the 


ALKALI-EARTH  METALS:    CALCIUM  453 

carbon  dioxide  which  will  be  in  equilibrium  with  a  mixture  of 
calcium  oxide,  CaO,  and  calcium  carbonate  is  as  follows  : 

Temperature         600°          700°          800°          898°       950° 
Pressure  of  CO2  2.35  mm.  25.3  mm.  168  mm.  760  mm.  1490  mm. 

It  will  be  seen  from  the  table  that  the  temperature  of  the  lime 
kiln  must  be  above  900°  for  the  carbon  dioxide  to  escape  rapidly 
from  the  interior  of  a  piece  of  limestone.  When  calcium  car- 
bonate is  heated  in  an  open  crucible  in  the  laboratory  under 
conditions  such  that  the  carbon  dioxide  is  constantly  displaced 
by  air,  the  partial  pressure  of  the  carbon  dioxide  may  become  very 
low  and  complete  decomposition  could  be  secured  at  a  tempera- 
ture of  750°  or  below. 

In  the  transformation  to  lime,  pieces  of  limestone  retain  their 
shape  but  shrink  somewhat  in  size.  If  water  is  added  to  the 
lime,  it  combines  with  it,  evolving  a  very  considerable  amount  of 
heat  —  15,540  small  calories  per  gram  molecule.  At  the  same 
time  the  calcium  hydroxide,  Ca(OH)2,  which  is  formed,  swells 
and  falls  to  a  loose  powder.  The  process  is  called  "  slaking."  If 
lime  is  exposed  to  the  air,  it  slowly  absorbs  water  and  falls  to  a 
powder,  and  the  hydroxide  also  absorbs  carbon  dioxide  and  is 
converted  back  to  the  carbonate.  Because  of  the  latter  change, 
"  air-slaked  lime  "  is  usually  worthless  for  the  preparation  of 
mortar. 

*  Dissociation  of  Calcium  Carbonate  and  the  Phase  Rule. 
In  the  system  produced  by  the  partial  dissociation  of  calcium 
carbonate  there  are  three  phases  (CaCO3,  CaO  and  C(>2,  and  only 
two  components  (CaO  and  CO2),  the  calcium  carbonate  formed 
by  the  union  of  the  other  two  not  being  considered  as  a  separate 
component.  According  to  the  phase  rule  (p.  107)  a  system  hav- 
ing two  components  and  three  phases  has  only  one  degree  of  free- 
dom and  is  univariant,  just  as  the  system  water  vapor  —  water 
has  one  component  and  two  phases  and  is  also  univariant. 
Accordingly,  in  the  system  calcium  carbonate  —  calcium  oxide 
—  carbon  dioxide,  if  the  temperature  changes,  the  pressure 
must  change  also ;  and  for  every  temperature  there  is  a  fixed 


454  A  TEXTBOOK  OF  CHEMISTRY 

dissociation  pressure,  just  as  for  every  temperature  there  is  a 
fixed  vapor  pressure  for  the  system  water  vapor  —  water. 

Mortar  is  prepared  by  mixing  slaked  lime  with  water  and  sharp 
sand,  which  has  not  been  rounded  by  long  action  of  waves. 
Where  used  between  bricks  the  moisture  is  partly  absorbed  by 
the  bricks  and  partly  dries  out  in  the  air.  This  is  followed  by 
the  action  of  the  carbon  dioxide  of  the  air,  which  slowly  changes 
the  hydroxide  to  calcium  carbonate  : 

Ca(OH)2  +  CO2  =  CaCO3  +  H2O 

The  carbonate  crystallizes  as  it  forms  and  adheres  strongly  to 
the  particles  of  sand,  binding  them  together.  In  plastered 
rooms  the  liberation  of  moisture  by  the  reaction  keeps  the  air 
of  the  rooms  moist  for  some  days  or  weeks. 

Cement.  For  the  manufacture  of  cement  a  clayey  limestone 
is  sometimes  used,  but,  as  limestones  having  the  proper  composi- 
tion are  rare,  artificial  mixtures  of  finely  powdered  limestone  and 
a  fine  clay  or  shale,  rich  in  silica,  are  usually  employed.  The 
slag  from  blast  furnaces  is  also  extensively  used.  The  materials 
are  finely  ground  and  fed  into  the  top  of  a  long,  slanting,  slowly 
rotating,  tubular  furnace  which  is  heated  by  means  of  powdered 
coal  blown  in  at  the  lower  end.  A  temperature  of  1500°-! 600° 
is  reached  in  the  hottest  parts  of  the  furnace.  The  carbon 
dioxide  is  completely  expelled  and  a  "  clinker  "  composed  of 
calcium  silicate  and  calcium  aluminate  with  an  excess  of  calcium 
oxide  is  formed.  This  clinker  is  finely  ground  and  mixed  with 
a  small  amount  of  plaster  of  Paris.  It  should  contain  only  a 
very  small  per  cent  of  magnesium.  The  composition  of  the 
finished  cement  is  as  follows : 

Loss  on  ignition  0-2  per  cent 

Silica,  SiO2  15-20  per  cent 

Alumina,  A12O3  3-8  per  cent 

Ferric  oxide,  Fe2Os  3-6  per  cent 

Lime,  CaO  58-64  per  cent 

Magnesium,  MgO  0-4  per  cent 

Potash  and  soda,  K2O,  Na2O  0-2  per  cent 

Sulfur  trioxide,  SOs  0-2  per  cent 


CALCIUM   CHLORIDE  455 

When  the  cement  is  mixed  with  water  it  slowly  combines  with 
it,  forming  partly  crystals  of  calcium  hydroxide,  Ca(OH)2,  partly, 
probably,  hydrated  silicates  and  aluminates  of  calcium  which 
"  set  "  to  a  firm,  strong  mass.  Sand  and  sometimes  other 
materials  are  usually  added  to  increase  the  volume.  As  the  set- 
ting results  from  the  action  of  water  alone,  it  will  take  place  under 
water  and  hence  the  material  is  often  called  hydraulic  cement. 

Calcium  Chloride,  CaCl2,  may  be  prepared  by  dissolving  cal- 
cium carbonate  in  hydrochloric  acid,  evaporating  the  solution 
and  drying  the  residue,  finally  at  a  temperature  of  260°  or  above. 
Prepared  in  this  manner  it  forms  a  porous,  extremely  hygro- 
scopic and  deliquescent  mass,  which  is  much  used  for  drying 
gases.  The  anhydrous  salt  melts  at  about  800°.  With  water 
it  forms  a  series  of  hydrates.  The  one  which  is  stable  at  ordinary 
temperatures  and  which  may  be  crystallized  by  cooling  very 
concentrated  solutions  is  CaCl2.6  H2O.  Considerable  heat  is 
evolved  when  the  anhydrous  chloride  dissolves  in  water,  but  a 
mixture  of  the  hydrate  with  a  little  less  than  its  weight  of  snow 
will  give  a  lowering  of  the  temperature  to  —  55°. 

Even  the  hydrate,  CaCl2.H2O,  has  an  appreciable  vapor  pres- 
sure at  ordinary  temperatures,  and  gases  cannot  be  fully  dried 
by  its  use.  At  15°  it  leaves  1.0  milligram  and  at  30°  it  leaves 
3.3  milligrams  of  water  in  liter  of  a  gas  which  has  been  passed 
over  it. 

Calcium  chloride  is  obtained  in  large  quantities  as  a  by-product 
in  the  ammonia-soda  process  (p.  412).  Because  of  the  low 
freezing  point,  solutions  of  calcium  chloride  are  used  in  refrigerat- 
ing machines  to  surround  the  cans  in  which  water  is  frozen  to 
artificial  ice.  They  have  also  been  used  for  sprinkling  roads. 

/OC1 
Chloride  of  Lime,"  Ca^         ,  is  prepared  commercially  by 

XC1 
spreading  slaked  lime  on  the  floor  of  a  room  and  filling  the  room 

with  chlorine  gas : 

OC1 

Ca(OH)2  +  C12  =  Ca/     '  +  H20 
XC1 


456  A  TEXTBOOK  OF  CHEMISTRY 

Treatment  with  a  dilute  acid  liberates  hydrochloric  and  hypo- 
chlorous  acids  and  these,  in  turn,  react  to  form  free  chlorine 
(p.  126).  The  compound  is  used  for  bleaching  (p.  125)  and  for 
sterilizing  water  (p.  83).  Good  bleaching  powder  should 
contain  35-37  per  cent  of  "  available  "  chlorine.  It  deteriorates, 
partly  by  loss  of  hypochlorous  acid  through  the  action  of  the 
carbon  dioxide  of  the  air,  partly  by  a  slow  transformation  of 
the  hypochlorite  to  the  chlorate. 

*  Calcium  Chlorate,  Ca(ClOs)2,  is  prepared  by  passing  chlorine 
into  "  milk  of  lime,"  a  mixture  of  calcium  hydroxide  and  water. 
The  hypochlorite  formed  at  first  changes  to  chlorate  by  auto- 
oxidation.  The  solution  is  used  to  mix  with  a  solution  of  potas- 
sium chloride,  KC1,  for  the  preparation  of  potassium  chlorate, 
KC1O3.  (Why?) 

Calcium  Fluoride,  CaF2,  is  found  as  the  mineral  fluorite,  which 
crystallizes  in  cubes  or  octahedra.  It  is  nearly  insoluble  in  water, 
and  for  this  reason  natural  waters  never  contain  more  than 
minute  traces  of  fluorine.  Calcium  fluoride  melts  at  1330°. 
It  is  used  as  a  flux  in  casting  iron  and  in  other  metallurgical  opera- 
tions and  is  the  source  from  which  hydrofluoric  acid  and  all  other 
compounds  of  fluorine  are  obtained. 

Calcium  Sulfide,  CaS,  is  formed  by  the  reduction  of  calcium 
sulfate,  CaSO4,  by  heating  it  with  coal  or  charcoal  in  the  Leblanc 
Soda  Process  (p.  411).  It  does  not  seem  to  be  appreciably 
soluble  in  water,  but  is  slowly  hydrolyzed  to  the  hydrosulfide, 
Ca(SH)2,  which  dissolves : 

2  CaS  +  2  HOH  =  Ca(SH)2  +  Ca(OH)2 

The  hydrolysis  is  sufficiently  slow  in  an  alkaline  solution  so 
that  the  sodium  carbonate  of  the  Leblanc  process  may  be  dis- 
solved, leaving  nearly  all  of  the  calcium  sulfide  as  an  insoluble 
residue. 

In  the  earlier  manufacture  the  calcium  sulfide  was  discarded 
as  a  waste  product,  but  the  slow  hydrolysis  of  the  material  on 
exposure  to  the  weather  led  to  the  contamination  of  streams  in 
England  and  resulted  in  stringent  legislation  requiring  the 


PLASTER   OF  PARIS  457 

manufacturers  to  take  care  of  their  waste  products  in  such  a 
manner  that  they  should  not  become  a  nuisance  to  others. 
This  and  the  competition  of  the  ammonia-soda  process,  com- 
pelling the  manufacturer  to  practice  every  possible  economy, 
led  to  the  invention  of  the  Chance  process  for  the  recovery  of 
the  sulfur.  By  exposing  the  moist  calcium  sulfide  to  the  action 
of  carbon  dioxide,  the  calcium  is  converted  to  the  carbonate  and 
hydrogen  sulfide  is  liberated : 

Ca(SH)2  +  H2C03  ^±  CaC03  +  2  H2S 

The  fact  that  the  ionization  constant  of  carbonic  acid  is  con- 
siderably greater  than  that  of  hydrogen  sulfide  and  also  the 
insolubility  of  the  calcium  carbonate,  both  aid  in  shifting  the 
equilibrium  of  the  reaction  to  the  right. 

By  burning  the  hydrogen  sulfide  with  a  limited  supply  of  air 
it  is  possible  to  convert  the  sulfur  almost  completely  to  the  free 
state. 

*  Acid  Calcium  Sulfite,  CaH2(SO3)2,  is  prepared  by  burning 
sulfur  and  passing  the  sulfur  dioxide,  SO2,  formed  into  milk  of 
lime.  The  solution  is  used  to  dissolve  and  remove  lignin  from 
the  fiber  in  the  manufacture  of  paper  from  wood. 

Calcium  Sulfate,  CaSO4.2  H2O,  Plaster  of  Paris.  The  min- 
eral gypsum,  which  has  the  composition  CaSO4.2  H2O,  is  found 
in  nature  in  clear,  transparent,  monoclinic  crystals  called  selenite 
and  in  a  white,  opaque  form  called  alabaster,  also  in  large  quan- 
tities in  a  form  suitable  for  making  plaster  of  Paris.  Calcium 
sulfate  is  also  sometimes  found  in  nature  in  the  anhydrous  form 
in  the  mineral  anhydrite.  If  gypsum  is  treated  for  some  time 
at  130°-160°,  it  loses  three  fourths  of  the  water  which  it  contains 
and  is  converted  into  a  compound  having  the  composition 
2  CaSO4.H2O  and  known  commercially  as  plaster  of  Paris. 
When  this  is  mixed  with  a  small  amount  of  water  to  a  creamy 
consistency,  it  can  be  filled  into  a  mold ;  but  after  standing  a  short 
time  part  of  the  water  unites  chemically  with  the  salt  to  form 
gypsum,  CaSO4.2  H2O,  which  "  sets  "  to  a  solid  mass.  If  the 
plaster  is  heated  to  too  high  a  temperature  or  for  too  long  a 


458  A  TEXTBOOK  OF  CHEMISTRY 

time  in  driving  out  the  water,  it  becomes  "  dead  burnt  "  and  will 
then  combine  with  water  only  very  slowly,  and  it  is  worthless 
for  the  ordinary  uses  of  plaster  of  Paris.  It  seems  probable  that 
ordinary  plaster  of  Paris  retains  a  small  amount  of  the  unchanged 
dihydrate  and  that  the  molecules  of  this  furnish  the  starting 
point  for  the  crystallization  in  setting.  Calcium  sulfate  is 
difficultly  soluble,  about  two  grams  dissolving  in  a  liter  of  water. 
It  is  much  less  soluble  in  alcohol. 

*  Plaster  of  Paris  and  the  Phase  Rule.  In  the  preceding 
paragraph  two  hydrates  of  calcium  sulfate,  CaSO4.2  H2O  and 
2  CaSO4.H2O,  have  been  mentioned  and  also  a  natural  anhy- 
dride, the  mineral  anhydrite.  Another  anhydride,  which  is 
more  easily  soluble  than  the  natural  anhydrite,  is  also  known  and 
is  called  the  "  soluble  anhydrite."  In  speaking  of  the  vapor 
pressure  of  hydrates  (p.  82)  it  has  been  implied  that  each 
hydrate  has  a  characteristic  vapor  pressure.  This  is  strictly 
true  only  in  case  the  loss  of  water  leads  to  the  formation  of  only 
one  compound,  either  a  lower  hydrate  or  an  anhydride.  The 
majority  of  salts  lose  water  in  this  manner;  but  gypsum, 
CaSO4.2  H2O,  may  lose  water  in  such  a  manner  as  to  form  either 
of  the  three  substances,  natural  anhydrite,  soluble  anhydrite  or 
plaster  of  Paris,  2  CaSO4.H2O.  The  vapor  pressure  of  the 
gypsum  will  depend  on  which  of  the  three  substances  is  formed, 
as  is  seen  in  the  table  on  the  opposite  page. 

If  these  values  are  plotted  and  the  vapor  pressure  curves 
extended,  it  is  found  (Fig.  101)  that  the  curve  for  the  system 
CaSO4.2  H2O,  CaSO4  (natural  anhydrite)  cuts  the  curve  for 
water  vapor  at  66°  and  the  other  two  curves  cut  it  at  89°  and 
107°.  As  there  are  only  two  components  (H2O  and  CaSO4), 
each  of  these  temperatures  represents  a  quadruple  point  (p.  78) 
where  the  sysem  is  invariant.  The  four  phases  at  66°  are: 
vapor,  solution,  CaSO4.2  H2O  and  CaSO4  (natural  anhydrite). 
Any  change  in  temperature  or  in  pressure  will  cause  the  disap- 
pearance of  one  of  the  phases.  If  the  pressure  is  decreased, 
water  will  evaporate  till  only  the  natural  anhydrite  is  left.  If 
the  pressure  is  increased,  the  vapor  phase  will  disappear.  If  the 


CALCIUM  SULFATE:    PHASE  RULE 


459 


VAPOR  PRESSURE  IN  MILLIMETERS  OF  MERCURY  FOR  SYSTEMS 
CONTAINING  GYPSUM 


t° 

PURE 
WATER 

SYSTEM 
CaSO4.2HzO 
AND  NATURAL 
ANHYDRITE 

SYSTEM 
CaSO4.2  H2O 
AND  SOLUBLE 
ANHYDRITE 

SYSTEM 
CaSO4.2  H2O 
AND  PLASTER  OF 
PARIS 
2  CaS04.H2O 

15 

12.7 

8.43 

7 

4.21 

20 

17.4 

12.2 

10.7 

6.24 

30 

31.5 

24 

19.4 

12.7 

40 

54.9 

45.4 

34 

26.3 

50 

149 

143 

108 

91.4 

65 

187 

140 

122 

70 

233 

185 

161 

80 

355 



314 

272 

90 

526 

446 

100 

760 





711 

105 

906 

888 

110 

1075 

1000  m  m 


900, 


800 


700 


600 


500 


400 


300 


60 


70° 


80° 


90°     100° 

Fig.  101 


110° 


460  A  TEXTBOOK  OF  CHEMISTRY 

temperature  is  increased,  all  of  the  gypsum,  CaSC>4.2  H2O, 
will  be  changed  to  anhydrite  because  the  vapor  pressure  of 
gypsum  would  be  greater  than  that  of  anhydrite  above  this 
temperature.  The  change  might,  it  is  true,  take  a  long  time,  but 
there  would  be  no  stable  equilibrium  till  the  change  was  com- 
plete. If  the  temperature  was  lowered,  all  of  the  anhydrite 
would  be  converted  into  gypsum,  since  the  vapor  pressure  of 
water  is  greater  than  that  of  gypsum  at  temperatures  below  66°. 

A  consideration  of  the  curves  will  enable  one  to  predict  whether 
gypsum  or  an  anhydrite  will  crystallize  from  a  salt  solution,  if 
we  know  the  vapor  pressure  of  the  latter.  Thus  a  solution  of 
sodium  chloride  has  a  vapor  pressure  greater  than  12.2  mm.  at 
20°.  From  such  a  solution  gypsum  will  crystallize.  If  mag- 
nesium chloride  or  calcium  chloride  is  added  till  the  vapor  pres- 
sure is  less  than  12.2  mm.  at  20°,  anhydrite  will  crystallize  from 
the  solution,  because  if  gypsum  and  such  a  solution  were  placed 
side  by  side  in  a  confined  space  water  would  escape  from  the 
gypsum  and  condense  in  the  solution.  These  predictions  of  the 
theory  agree  with  facts  observed  by  geologists  about  the  condi- 
tions under  which  gypsum  and  anhydrite  are  found  in  nature. 

*  Calcium  Nitrate,  Ca(NO3)2-4  H2O,  is  readily  prepared  by 
dissolving  the  carbonate  in  nitric  acid.  It  is  so  manufactured 
for  fertilizers  and  other  uses  on  a  considerable  scale  in  Norway, 
by  absorbing  the  oxides  of  nitrogen,  formed  by  the  electric  arc 
from  air,  in  milk  of  lime.  The  anhydrous  salt  is  sometimes  used 
to  dry  oxides  of  nitrogen  or  other  gases  for  which  calcium  chloride 
cannot  well  be  used. 

Calcium  Phosphates.  Normal  calcium  phosphate,  Ca3(PO4)2, 
is  an  important  constituent  of  bones.  It  is  also  found  as  a  min- 
eral phosphate  in  deposits  in  the  southeastern  part  of  the 
United  States  from  North  Carolina  and  Tennessee  to  Florida. 
Both  bones  and  the  mineral  phosphates  are  extensively  used  as 
fertilizers,  phosphorus  being  an  element  which  is  essential  for 
the  growth  of  crops  and  which  is  found  in  only  limited  amounts 
in  some  soils.  In  order  to  render  the  phosphate  more  easily 
soluble  and  available  for  the  growth  of  plants,  the  powdered 


CALCIUM   PHOSPHATES  461 

mineral  phosphate  is  often  treated  with  sulfuric  acid  to  convert 
it  into  monocalcium  phosphate,  CaH4(PO4)2.  The  mixture  of 
calcium  sulfate  and  acid  calcium  phosphate  is  designated,  com- 
mercially, as  a  "  superphosphate."  A  slightly  diluted  acid  is 
used  and  this  "superphosphate  "  contains  both  salts  as  hydrates  : 

Ca3(PO4)2  +  2  H2SO4  +  6  H2O 

=  Ca(H2PO4)2.2  H2O  +  2  (CaSO4.2  H2O) 

*  Solubility  of  Calcium  Phosphates.  Even  the  monocalcium 
phosphate,  Ca(H2PO4)2,  is  only  slightly  soluble,  about  5  grams 
dissolving  in  a  liter  of  water.  If  a  larger  amount  of  the  salt  is 
added,  the  reaction  : 

2  H2PO4-  ^±  HPO4~  +  H3PO4 

causes  the  number  of  monohydrophosphate  ions,  HPO4  ,  and 
of  calcium  ions  in  the  solution  to  exceed  the  solubility  product  for 
the  reaction, 


and  there  will  be  formed  a  precipitate  of  the  very  difficultly 
soluble  dicalcium  phosphate,  CaHPO4,  while  the  solution  will 
contain  more  phosphoric  acid  than  will  correspond  to  the  mono- 
calcium  phosphate,  Ca(H2PO4)2.  In  the  valuation  of  fertilizers 
it  is  customary  to  distinguish  three  forms  of  phosphoric  acid, 
"  water-soluble  phosphoric  acid,"  "  citrate-soluble  phosphoric 
acid  "  (phosphoric  acid  insoluble  in  water  but  soluble  in  a  neutral 
solution  of  ammonium  citrate  of  sp.  gr.  1.09)  and  "  insoluble 
phosphoric  acid."  The  last  is  supposed  to  be  in  the  form  of 
tricalcium  phosphate,  Caa(PO4)2.  From  what  has  been  said 
above  it  is  evident  that  more  phosphoric  acid  will  pass  into 
solution  if  a  dicalcium  phosphate  is  treated  with  successive  small 
portions  of  water  than  if  it  is  treated  at  once  with  a  large  quan- 
tity of  water.  The  amount  of  the  "  citrate  soluble  phosphoric 
acid  "  also  depends  on  the  exact  conditions  of  the  determination. 
Both  determinations  are  to  be  considered  as  in  a  considerable 
measure  conventional,  and  it  is  quite  certain  that  they  do  not 
furnish  an  accurate  measure  of  the  availability  of  the  phosphorus 


462  A  TEXTBOOK  OF  CHEMISTRY 

for  the  growth  of  plants.  Finely  ground  phosphate  rock  also 
furnishes  phosphorus  which  can  be  slowly  absorbed  by  plants, 
and  it  is  at  least  a  question  whether  the  acid  phosphate  is  any 
better  than  the  raw  ground  phosphate  for  maintaining  the 
fertility  of  land  during  a  series  of  years. 

The  addition  of  even  a  weak  acid,  such  as  acetic  acid,  to  cal- 
cium phosphate  will  cause  it  to  pass  into  solution.  This  is 
because  phosphoric  acid  is  so  weak  an  acid  that  in  the  presence 
of  even  comparatively  few  hydrogen  ions,  H+,  the  phosphoric 
acid  must  be  almost  completely  either  in  the  form  of  the  un-ion- 
ized  acid,  H3PO4,  or  of  the  dihydrogen  phosphate  ions,  H2PO4-. 
Under  these  conditions  there  can  not  be  enough  of  the  mono- 
hydrogen  phosphate  ions  present  to  reach  the  solubility  product  : 

Ca++  X  HPO4~  =  K 

which  would  cause  the  precipitation  of  the  dicalcium  phosphate. 
As  has  been  explained  above,  it  is  this  precipitation  which 
causes  the  apparent  difficult  solubility  of  monocalcium  phos- 
phate. 

Calcium  Carbide,  CaC2.  By  heating  a  mixture  of  lime,  CaO, 
and  coke  in  a  revolving  electric  furnace,  calcium  carbide  is 
formed  : 


3  C  =  CaC;2  +  2  CO 

The  carbide  is  easily  hydrolyzed  by  water  with  the  formation 
of  calcium  hydroxide  and  acetylene,  C2H2  (p.  293),  and  is  manu- 
factured chiefly  for  that  use. 

Calcium  Cyanamide,  CaCN2.  By  heating  calcium  carbide 
in  a  current  of  nitrogen  at  1000°  it  is  transformed  into  calcium 
cyanamide,  Ca=N  —  C=N  : 

CaC2  +  2  N  =  CaCN2  +  C 

Calcium  cyanamide  is  hydrolyzed  by  water  to  ammonia  and 
calcium  carbonate  : 

CaCN2  +  3  H20  =  CaCO3  +  2  NH3 

As  these  reactions  furnish  a  means  of  transforming  the  nitrogen 
of  the  air  into  a  form  which  is  available  for  plant  growth,  calcium. 


HARD  WATERS  463 

cyanamide  is  now  manufactured  in  considerable  quantities  for 
use  as  a  fertilizer.  For  this  use  it  is  often  called  "  lime-nitro- 
gen," or  in  German  "  Kalk-Stickstoff." 

Calcium  Carbonate,  CaCO3,  is  the  most  abundant  compound 
of  calcium.  The  various  forms  have  been  already  mentioned, 
also  its  conduct  when  heated. 

Hard  Waters.  One  liter  of  pure  water  at  ordinary  tempera- 
tures will  dissolve  only  about  12  milligrams  of  calcium  carbonate, 
CaCOs,  but  water  saturated  with  carbon  dioxide  will  dissolve 
nearly  100  times  as  much,  or  more  than  a  gram  in  one  liter. 
The  solution  contains  the  acid  carbonate,  Ca(HCO3)2,  frequently 
called  the  bicarbonate.  As  natural  waters  always  contain  some 
carbon  dioxide  absorbed  from  the  air  and  usually  acquire  much 
more  from  decaying  vegetable  matter  in  the  soil,  all  such  waters 
which  have  come  in  contact  with  a  soil  containing  calcium  car- 
bonate hold  more  or  less  of  the  calcium  bicarbonate  in  solution. 

The  properties  of  such  a  solution  can  be  easily  illustrated  by 
passing  carbon  dioxide  through  a  solution  of  limewater,  slightly 
diluted.  Calcium  carbonate  will  be  precipitated  at  first,  but  on 
continuing  the  current  of  the  gas  it  will  pass  again  into  solution. 
On  boiling  the  solution  the  acid  carbonate  dissociates  into 
calcium  carbonate,  carbon  dioxide  and  water.  The  carbon 
dioxide  escapes  with  the  steam  and  the  calcium  carbonate  is 
precipitated.  Natural  waters  of  this  type  are  said  to  have 
"  temporary  hardness,"  since  the  hardness  is  nearly  all  removed 
by  boiling.  The  designation  "  hardness  "  refers  to  the  effect 
of  hard  water  in  precipitating  an  insoluble  calcium  salt  when 
soap  is  used  with  it.  The  water  continues  to  have  a  harsh  feel- 
ing to  the  skin  until  enough  soap  has  been  used  to  complete  the 
precipitation  of  the  calcium  and  magnesium  salts  which  are 
present. 

Waters  containing  calcium  sulfate  in  solution  will  not  deposit 
the  salt  on  short  boiling  and  are  said  to  have  "  permanent 
hardness."  Such  waters  deposit  the  sulfate  on  concentration 
of  the  water,  as  is  done  in  a  steam  boiler,  and  the  decreased 
solubility  of  the  calcium  sulfate  at  a  high  temperature  increases 


464  A  TEXTBOOK  OF  CHEMISTRY 

the  amount  of  scale  formed.     The  scale  from  such  a  water  is 
especially  coherent  and  troublesome. 

If  milk  of  lime,  Ca(OH)2,  is  added  to  water  containing  calcium 
bicarbonate  in  just  the  right  proportion,  nearly  all  of  the  calcium 
carbonate  in  the  water  and  also  the  calcium  of  the  milk  of  lime 
which  is  added  will  be  precipitated.  This  is  known  as  Clark's 
process  of  softening  water  : 

Ca(HCO3)2  +  Ca(OH)2  =  2  CaCO3  +  2  H2O 

To  remove  the  calcium  of  calcium  sulfate,  sodium  carbonate, 
trisodium  phosphate,  sodium  fluoride  or  some  other  salt  which 
will  precipitate  the  calcium  must  be  used. 

*  Determination  of  Free  and  Combined  Carbonic  Acid  in 
Natural  Waters.  It  has  been  shown  (p.  389)  that  phenol- 
phthalein  is  a  suitable  indicator  for  weak  acids  ;  and  methyl 
red,  for  weak  bases.  If  the  acid  is  very  weak  indeed,  the 
hydrolysis  of  the  normal  salt  may  cause  the  end  point  with 
phenolphthalein  to  appear  before  the  acid  is  fully  neutralized. 
In  the  case  of  carbonic  acid,  H2COs,  when  the  point  correspond- 
ing to  the  formation  of  NaHCOs  is  passed,  the  hydrolysis  of  the 
sodium  carbonate, 


Na+  +  Na+  +  CO3=  +  H+ 

=  Na+  +  Na+  +  HC(V  +  OH~ 

causes  the  concentration  of  hydroxide  ions  to  exceed  the  end 
point  for  phenolphthalein.  The  amount  of  alkali  required  to 
give  a  pink  color  with  phenolphthalein  will,  therefore,  indicate 
the  amount  of  free  carbonic  acid,  H2COs,  in  the  solution. 

On  the  other  hand,  if  a  strong  acid  is  added  to  a  solution  of  a 
carbonate  or  bicarbonate,  so  long  as  any  of  either  remains  in 
solution  the  concentration  of  the  hydrogen  ions  cannot  exceed 
the  concentration  in  a  solution  of  carbonic  acid,  H2CO3.  When 
enough  acid  has  been  added  to  decompose  all  of  the  carbonates 
and  bicarbonates  present,  any  further  addition  of  acid  will 
carry  the  concentration  of  the  hydrogen  ions  past  the  end  point 
for  methyl  orange  or  methyl  red.  This  makes  it  possible  to 


CALCIUM  OXALATE  465 

determine  the  amount  of  free  carbonic  acid  and  also  of  carbonates 
and  bicarbonates  present,  by  titrating  first  with  alkali,  using 
phenolphthalein  as  an  indicator  and  then  with  an  acid,  using 
methyl  orange  or  methyl  red. 

*  Calcium  Acetate,  Ca(C 2X1302)2,  is  prepared  commercially  by 
neutralizing  the  distillate  obtained  by  the  destructive  distilla- 
tion of  wood.  It  is  used  for  the  manufacture  of  glacial  acetic 
acid,  HC2H3O2,  and  of  acetone,  CH3COCH3. 

Calcium  Oxalate,  CaC2O4.H2O.  When  a  solution  of  ammo- 
nium oxalate,  (NH4)2C2O4,  is  added  to  a  solution  containing  a 
soluble  salt  of  calcium,  calcium  oxalate  is  precipitated  as  a  fine, 
crystalline  powder.  One  liter  of  water  dissolves  only  5.6 
milligrams  of  the  salt,  and  it  is  still  less  soluble  in  a  solution  con- 
taining ammonium  oxalate.  For  this  reason  it  is  often  used  for 
the  quantitative  determination  of  calcium.  Its  value  for  this 
purpose  is  greatly  enhanced  by  the  fact  that  magnesium  oxalate, 
MgC2O4,  is  much  more  easily  soluble  —  300  milligrams  in  one 
liter  of  water. 

In  strong  acids,  as  hydrochloric  or  nitric  acid,  calcium  oxalate 
dissolves.  Oxalic  acid  is  only  moderately  ionized  in  solutions  of 
medium  concentration  —  50  per  cent  to  H+  and  HC2O4~  in 
tenth  normal  solution.  The  ionization  to  H+,  H+  and  €204 
must  be  very  much  less.  As  the  presence  of  the  hydrogen  ions 
of  a  highly  ionized  acid,  such  as  hydrochloric  acid,  shifts  the 
equilibrium  of  the  reaction  : 

H2C204  ^±  H+  +  H+  +  C204- 

to  the  left,  the  concentration  of  the  oxalate  ions,  C2O4  ,  in  such 
a  solution  cannot  be  great  enough  for  the  solubility  product 
C2O4  X  Ca++  to  reach  the  point  of  precipitation  for  calcium 
oxalate.  The  salt  will,  therefore,  dissolve  in  such  a  solution. 
In  a  solution  of  a  weak  acid  such  as  acetic  acid,  HC2H3O2,  the 
number  of  hydrogen  ions  is  so  small  that  they  produce  only  a 
slight  effect  on  the  ionization  of  oxalic  acid,  which  is  a  very 
much  stronger  acid  than  acetic  acid.  Calcium  oxalate  may  be 
precipitated,  for  this  reason,  from  solutions  containing  acetic 


466  A  TEXTBOOK  OF   CHEMISTRY 

acid,  though  the  salt  is  more  soluble  in  dilute  acetic  acid  than 
in  pure  water.  Owing  to  the  extremely  low  ionization  constant 

f  or :  H2PO4-+  ^±  H+  +  HPO4" 

calcium  phosphate  is  not  precipitated  from  solutions  containing 
even  the  very  weak  acid,  acetic  acid.  The  addition  of  ammo- 
nium oxalate  to  a  solution  of  calcium  phosphate  in  dilute  acetic 
acid  will,  accordingly,  cause  the  precipitation  of  nearly  all  of  the 
calcium  in  the  form  of  the  oxalate,  and  even  a  solution  of  calcium 
sulfate  will  precipitate  oxalic  acid,  but  not  phosphoric  acid,  from 
a  solution  of  an  oxalate  containing  acetic  acid. 

Calcium  oxalate  loses  carbon  monoxide  and  is  converted 
into  calcium  carbonate  by  gentle  ignition : 

CaC2O4  =  CaCO3  +  CO 

Calcium  Silicate,  CaSiO3,  is  found  in  nature  as  the  mineral 
wollastonite.  It  may  be  prepared  by  fusing  a  mixture  of  quartz 
or  sand  with  calcium  carbonate  : 

CaCO3  +  Si02  =  CaSiO3  +  CO2 

Calcium  silicate  is  a  constituent  of  a  very  large  proportion  of 
the  most  common  natural  silicates,  such  as  pyroxene,  amphibole, 
garnet  and  the  zeolites. 

Glass.  By  melting  together  a  mixture  of  calcium  carbonate, 
CaCO3,  sodium  carbonate,  Na2CO3,  and  a  pure  quartz  sand, 
SiO2,  in  proper  proportions,  a  silicate  of  calcium  and  sodium  is 
obtained,  which  does  not  readily  crystallize  on  cooling,  but  which 
passes  through  a  stage  in  which  it  becomes  more  and  more  viscous 
and  finally  solidifies  to  a  transparent,  homogeneous  mass.  The 
sodium  may  be  partly  or  completely  replaced  by  potassium,  the 
calcium  may  be  replaced  by  lead,  and  part  of  the  silica  by  boric 
anhydride,  B2O3,  giving  glasses  suitable  for  special  uses.  In  all 
of  these  the  glass  is  to  be  considered  as  a  complex  mixture  of 
silicates  in  the  form  of  an  extremely  viscous,  supercooled  liquid. 
The  value  of  glass  depends  largely  upon  the  fact  that  as  an 
amorphous,  supercooled  liquid  it  is  still  viscous  but  sufficiently 


GLASS.    STRONTIUM  467 

plastic  so  that  it  can  be  worked  over  a  considerable  range  of 
temperature. 

Some  of  the  more  important  varieties  of  glass  are  as  follows : 
window  glass,  plate  glass  and  the  glass  of  ordinary  table  ware  are 
usually  a  silicate  of  calcium  and  sodium.  The  finer  grades  of 
such  glass  are  often  called  crown  glass.  Flint  glass  is  a  silicate 
of  lead  and  sodium  or  potassium  prepared  by  melting  together 
litharge,  PbO,  potassium  carbonate,  K2CO3,  and  silica,  SiO2. 
The  name  comes  from  the  former  use  of  crushed  flints  for  the 
silica.  Flint  glass  has  a  higher  index  of  refraction  but  also  a 
relatively  greater  dispersive  power  than  crown  glass,  and  the  two 
varieties  of  glass  are  used  together  for  achromatic  lenses  and  for 
direct  vision  spectroscopes.  It  melts  easily.  Strass  or  paste 
is  a  heavy  lead  glass  with  a  high  index  of  refraction,  used  in 
making  imitations  of  diamonds  and  other  precious  stones. 
Bohemian  glass  is  a  silicate  of  potassium  and  calcium  having  a 
high  melting  point.  It  is  also  less  soluble  in  water  than  ordinary 
glass  and  is  used  for  combustion  tubing  and  for  beakers  and 
flasks  used  in  the  laboratory.  It  is  often  called  hard  glass.  It 
has  been  largely  replaced  by  various  borosilicate  glasses. 
"  Jena  "  glass,  "  Resistanz  "  glass  and  "  Non-sol  "  glass  are 
borosilicates  containing  a  little  zinc.  They  are  much  less  sol- 
uble in  water  than  the  ordinary  glasses,  and  some  of  them  soften 
at  much  higher  temperatures.  These  properties  render  them 
suitable  for  special  uses  in  chemical  laboratories,  especially  for 
combustion  tubing,  for  beakers  and  flasks  for  use  in  quantitative 
analysis  and  for  test  tubes  for  bacteriological  cultures.  Durax 
glass  is  a  variety  especially  resistant  to  alkaline  solutions.  For 
the  manufacture  of  thermometers  several  borosilicate  glasses 
are  made  which  give  a  much  smaller  depression  of  the  zero  point 
after  use  at  high  temperatures  than  is  the  case  with  thermometers 
made  from  ordinary  glass.  Some  of  these  glasses  also  give 
thermometers  which  correspond  closely  with  the  hydrogen 
thermometer  at  high  temperatures. 

Strontium,  Sr,  87.63.  Occurrence.  As  the  most  important 
natural  compounds  of  calcium  are  the  carbonate  and  sulfate,  so 


468  A  TEXTBOOK  OF  CHEMISTRY 

strontium  and  barium  are  found  chiefly  as  carbonates  and 
sulfates.  The  sulfates  are  much  less  soluble  than  calcium  sul- 
fate.  Strontium  carbonate,  SrCOs,  is  called  strontianite,  and 
strontium  sulfate,  SrSO4,  celestite.  The  latter  name  is  given 
because  the  mineral  is  often  of  a  light  blue  color,  but  the  pure 
sulfate  is  white.  Water  dissolves  about  20  times  as  much  cal- 
cium sulfate  as  it  does  of  strontium  sulfate. 

*  Strontium  Hydroxide,  Sr(OH)2.8  H2O,  is  more  soluble  than 
calcium  hydroxide  and  is  much  more  easily  soluble  in  hot  than  in 
cold  water.  It  forms  a  difficultly  soluble  compound  with  cane 
sugar,  Ci2H22On,  and  is  sometimes  used  to  recover  sugar  from 
the  molasses  of  the  beet  sugar  manufacture. 

Strontium  Nitrate,  Sr(NO3)2,  is  used,  mixed  with  sulfur, 
charcoal  and  potassium  chlorate  or  nitrate  for  red  lights  in  fire- 
works. 

Barium,  Ba,  137.37.  Occurrence.  Barium  is  found  in  nature 
as  the  carbonate,  BaCOs,  called  witherite,  and  the  sulfate,  BaSC>4, 
called  barite.  The  latter  is  more  common. 

Barium  Oxide,  BaO.  Barium  carbonate  is  very  much  more 
stable  than  calcium  carbonate.  The  dissociation  pressure  of 
calcium  carbonate  reaches  atmospheric  pressure  at  898°,  but  that 
of  barium  carbonate  is  equal  to  one  atmosphere  at  1350°  (Finket- 
stein,  Ber.  39,  1588  (1906)).  Even  at  this  temperature  the 
decomposition,  at  first,  appears  to  give  a  basic  carbonate,  prob- 
ably BaCO3.BaO,  and  a  temperature  of  1450°  is  required  for 
the  decomposition  of  this  compound  at  atmospheric  pressure. 
It  is  impracticable,  therefore,  to  prepare  barium  oxide  by  the 
direct  decomposition  of  the  carbonate.  If  the  carbonate  is 
mixed  with  charcoal,  however,  the  carbon  dioxide  formed  by 
the  dissociation  of  the  barium  carbonate  will  at  once  react  with 
the  carbon  to  form  carbon  monoxide.  As  a  result  of  the  two 
equilibria:  BaCO3  ^±  BaO  +  CO2 

co2  +  c  :£  2  co 

the  constant  removal  of  the  carbon  dioxide  by  means  of  the 
second  reaction  makes  it  possible  to  prepare  barium  oxide  at  a 


ALKALI-EARTH  METALS:  BARIUM  469 

temperature  at  which  the  dissociation  pressure  is  small  and  the 
preparation  is  carried  out,  technically,  by  this  method. 

Barium  oxide  may  also  be  prepared  by  the  decomposition  of 
the  nitrate,  Ba(NOa)2,  which  occurs  at  a  much  lower  temperature 
than  that  for  the  decomposition  of  the  carbonate. 

Barium  oxide  combines  directly  with  water  to  form  the  hydrox- 
ide, Ba(OH)2,  the  reaction  being  accompanied  by  considerable 
evolution  of  heat.  It  is  used,  chiefly,  for  the  preparation  of  the 
peroxide,  BaO2.  It  is  also  used  in  the  laboratory  as  a  powerful 
dehydrating  agent,  as  in  the  preparation  of  absolute  alcohol. 

Barium  Peroxide,  BaO2.  The  dissociation  pressure  for  the 
reaction,  2  BaO  +  O2  ^±  2  BaO2,  is : 

Temperature         525°        670°        735°         775°          790° 
Pressure  in  mm.      20  80  260  510  670 

When  we  remember  that  the  partial  pressure  of  oxygen  in  the 
air  is  21  per  cent  of  760  mm.,  or  160  mm.,  it  is  evident  that  if 
air  is  passed  over  barium  oxide  at  a  temperature  of  670°,  oxygen 
will  be  absorbed  with  the  formation  of  the  peroxide.  On  the 
other  hand,  if  the  peroxide  is  heated  to  about  800°  under  atmos- 
pheric pressure,  the  peroxide  will  be  decomposed  with  evolution 
of  oxygen.  Again,  if  barium  oxide  at  800°  is  subjected  to  the 
action  of  air  under  a  pressure  of  about  five  atmospheres,  so  that 
the  partial  pressure  of  the  oxygen  is  a  little  more  than  one 
atmosphere,  oxygen  will  be  absorbed,  and  on  lowering  the  pres- 
sure to  atmospheric  pressure  oxygen  will  be  evolved  without 
any  change  in  temperature.  Both  methods  have  been  used  for 
the  technical  preparation  of  oxygen.  A  little  water  vapor  must 
be  present  to  catalyze  the  reaction,  and  the  carbon  dioxide  of 
the  air  must  be  carefully  removed. 

Barium  peroxide  is  also  prepared  and  used  for  the  manu- 
facture of  hydrogen  peroxide,  H2O2.  It  forms  a  hydrate, 
BaO2.8  H2O,  which  is  difficultly  soluble  and  which  is  formed  on 
treating  the  anhydrous  barium  peroxide  with  water.  It  is  also 
formed  by  precipitating  an  ice-cold  solution  of  barium  chloride, 
BaCl2,  with  a  cold  solution  of  sodium  peroxide,  Na2O2. 


470  A  TEXTBOOK  OF  CHEMISTRY 

Barium  Hydroxide,  Ba(OH)2.8  H^O,  is  much  more  easily 
soluble  in  water  than  calcium  or  strontium  hydroxides.  It  is 
also  much  more  easily  soluble  in  hot  than  in  cold  water  and  can 
be  easily  recrystallized  from  hot  water.  It  gives  a  precipitate  of 
barium  carbonate,  BaCOs,  with  carbon  dioxide,  and  is  a  very 
sensitive  reagent  for  the  qualitative  detection  or  quantitative 
determination  of  that  gas.  It  is  used  to  prepare  an  alkali 
solution  which  is  free  from  carbonate,  but  such  a  solution  must, 
of  course,  be  carefully  protected  from  the  carbon  dioxide  of  the 
air. 

Barium  Chloride,  BaCl2.2  H^O,  is  the  most  common  sol- 
uble salt  of  barium.  It  is  used  especially  for  the  detection  and 
quantitative  determination  of  sulfates. 

*  Barium  Nitrate,  Ba(NOa)2,  is  sometimes  used  for  the  detec- 
tion of  sulfates  in  solution  containing  silver  or  other  metals  which 
form  insoluble  chlorides. 

*  Barium  Sulfide,  BaS,  is  formed  "by  hefting  barium  sulfate, 
the  most  plentiful  natural  source  of  barium  compounds,  with 
charcoal :  g^  +  4  c  =  BaS  +  4  CO 

Barium  sulfide  dissolves  in  hydrochloric  acid  with  the  forma- 
tion of  the  chloride  or  in  nitric  acid  with  the  formation  of  the 
nitrate,  methods  which  were  formerly  used  for  the  preparation 
of  these  salts  from  the  insoluble  sulfate. 

Barium  Sulfate,  BaSO4.  The  mineral  barite  is  sufficiently 
abundant  to  form  a  cheap  source  for  barium  compounds.  It  is 
also  sometimes  used  in  a  finely  pulverized  form  as  an  adulterant 
for  white  lead.  It  is  crystalline  and  the  particles  are  much  more 
transparent  than  those  of  white  lead,  so  that  paints  containing 
it  have  less  covering  power  than  those  made  from  pure  lead 
compounds.  It  is  used  in  many  mixed  paints,  especially  in 
"  lithopone,"  which  is  a  mixture  of  barium  sulfate,  BaSO4,  and 
zinc  sulfide,  ZnS,  obtained  by  precipitating  a  solution  of  zinc 
sulfate  with  barium  sulfide,  BaS.  Lithopone  is  not  blackened 
by  hydrogen  sulfide  and  is  more  suitable  than  white  lead  for 
places  where  that  gas  is  liable  to  be  present. 


RADIUM  471 

Barium  sulfate  requires  about  400,000  parts  of  water  for  its 
solution  and  it  is  only  slightly  more  soluble  in  dilute  acids.  It 
is  very  much  used  for  this  reason  for  the  detection  and  estima- 
tion of  sulfates.  It  has  a  very  marked  tendency  to  form  solid 
solutions  or  mixed  crystals  with  other  sulfates,  barium  chloride 
or  other  compounds  which  may  be  present.  The  presence  of 
these  foreign  substances  may  lead  to  rather  serious  errors  in  the 
determination  of  sulfates,  unless  great  care  is  used  in  following 
proper  methods  of  manipulation. 

Barium  sulfate  is  appreciably  soluble  in  solutions  containing 
considerable  hydrochloric  acid  and  dissolves  rather  easily  in 
concentrated  sulfuric  acid. 

Flame  Colors  for  Calcium,  Strontium,  and  Barium.  Calcium 
compounds  impart  a  brick  red  color  to  the  Bunsen  flame,  stron- 
tium compounds  a  bright  red  and  barium  compounds  a  green 
color.  The  spectra  show  bright  lines  and  bands  which  are  easily 
distinguished,  even  f^fh  the  simplest  forms  of  spectroscopes. 

Radium  (Ra,  226.4).  About  1878  Sir  William  Crookes  dis- 
covered that  in  the  discharge  of  electricity  through  a  highly 
rarefied  gas  rays  are  shot  out  at  right  angles  to  the  surface  of  the 
cathode  and  produce  a  beautiful  green  fluorescence  at  the  point 
where  they  strike  the  containing  tube.  About  twenty  years 
later  it  was  shown  that  these  rays  consist  of  electrons  traveling 
with  velocities  approaching  that  of  light.  In  1895  Rontgen, 
partly  by  accident,  discovered  that  rays,  afterwards  called 
Rontgen  rays,  emanate  from  the  glass  at  the  point  of  fluores- 
cence. These  rays  penetrate  paper,  wood  and  some  other  ob- 
jects, which  are  opaque  to  ordinary  light,  and  affect  a  photo- 
graphic plate  placed  behind  such  screens.  The  rays  are,  however, 
intercepted  by  metals  or  substances  containing  compounds  of  the 
metals,  as,  for  instance,  by  bones.  The  opacity  of  various  sub- 
stances is  closely  proportional  to  their  density.  Very  shortly 
after,  in  1896,  Becquerel,  in  Paris,  discovered  that  minerals  con- 
taining uranium  have  the  property  of  emitting  penetrating  radi- 
ations resembling  the  Rontgen  rays  in  their  effect  on  a  photo- 
graphic plate  protected  from  ordinary  light  rays  by  a  screen  of 


472  A  TEXTBOOK  OF  CHEMISTRY 

black  paper.  Further  study  by  Madame  and  Monsieur  Curie 
led  to  the  discovery  in  uranium  ores  of  a  new  element,  radium. 
This  belongs  to  the  calcium-barium  family,  and  forms  a  sulfate 
which  is  much  less  soluble  than  barium  sulfate  and  which  is  made 
use  of  in  separating  radium  from  other  elements.  Radium  and 
its  compounds  exhibit  the  following  remarkable  properties : 

1.  It  affects  a  photographic  plate  through  black  paper  and 
will  cause  the  fluorescence  of  zinc  sulfide  and  some  other  com- 
pounds exposed  to  the  action  of  its  rays. 

2.  It  causes  the  ionization  of  air,  that  is,  the  separation  of  the 
molecules  of  the  air  into  positive  and  negative,  charged  particles, 
causing  the  air  to  become  a  conductor  of  electricity.     A  gold-leaf 
electroscope,  which  will  remain  charged  for  a  long  time  in  ordi- 
nary air,  is  rapidly  discharged  and  the  leaves  fall  together  in  air 
which  has  been  exposed  to  the  action  of  radium.     The  rate  of 
discharge  furnishes  a  quite  accurate  measure  of  the  quantity  of 
radioactive  substances  present,  and  the  measurement  of  this  rate 
is  the  most  important  method  used  in  the  study  of  radioactive 
elements. 

3.  Radium  continually  evolves  heat.     One  gram  of  the  ele- 
ment gives  out  132  small  calories  per  hour.     This  phenomenon 
is  independent  of  the  temperature  or  of  the  form  in  which  the 
element  is  combined  or  of  any  other  conditions  which  can  be 
controlled. 

Disintegration  of  Atoms.  In  1902-1903  Professor  Rutherford, 
then  at  McGill  University  in  Montreal,  published  a  series  of 
papers  in  which  he  proposed  the  hypothesis  that  the  atoms  of 
radioactive  elements  disintegrate  more  or  less  rapidly,  breaking 
down  with  the  formation  of  other  elements.  In  the  disintegra- 
tion, portions  of  the  atom  are  shot  out  from  it  with  tremendous 
velocity.  Some  sort  of  potential  or  kinetic  energy  within  the 
atom  is  liberated  in  this  manner  and  manifests  itself  as  heat 
energy,  and  this  explains  the  heat  evolved  by  radioactive  ele- 
ments. Incidentally  this  makes  it  probable  that  all  atoms  are 
complex  in  their  internal  structure  and  are  storehouses  of 
immense  quantities  of  energy.  The  particles  shot  out  by  the 


RADIUM:    DISINTEGRATION  THEORY  473 

disintegrating  atoms  seem  to  tear  apart  molecules  which  they 
strike,  separating  them  into  charged  particles  or  ions,  and  in  this 
manner  air  or  other  gases  which  are  exposed  to  the  action  of 
radioactive  substances  become  conductors  of  electricity  as  has 
been  mentioned  above. 

Rutherford's  disintegration  theory  was  very  strongly  sup- 
ported when  Soddy,  who  began  work  with  Rutherford,  demon- 
strated, while  working  with  Sir  William  Ramsay  in  London,  that 
helium  is  one  of  the  disintegration  products  of  radium.  The 
theory  is  now  accepted,  at  least  as  a  working  hypothesis,  by  all 
investigators  in  this  field.  A  few  chemists,  on  account  of  the 
disintegration  of  radium,  have  contended  that  it  is  not  properly 
called  an  element  but  should  be  classed  as  a  chemical  compound. 
But  radium  finds  its  place  in  the  periodic  system  and  has  all  of 
the  other  properties  which  usually  characterize  an  element. 
Moreover,  there  is  good  reason  to  believe  that  radium  is,  in  turn, 
a  disintegration  product  of  uranium.  If  radium  is  a  compound, 
uranium  is  a  compound  also.  It  seems  better,  therefore,  to 
revise  our  definition  of  an  element  and  accept  the  notion  that 
the  atoms  of  some  elements,  and  possibly  of  all,  may  disintegrate 
with  the  formation  of  other  atoms.  The  relation  between  an 
atom  and  its  disintegration  products  is  evidently  very  different 
from  the  relation  between  a  compound  and  the  elements  of  which 
it  is  composed. 

Nature  of  the  Radiations  from  Radioactive  Substances. 
Four  kinds  of  rays  have  been  distinguished  as  emanating  from 
radioactive  substances.  The  first  kind,  called  «-rays,  have  been 
identified  as  atoms  of  helium,  carrying  a  double  positive  charge. 
As  they  are  charged  particles  moving  with  a  high  velocity,  they 
are  slightly  deflected  by  a  strong  magnetic  field.  The  /8-rays 
appear  to  be  identical  with  the  cathode  rays  of  the  Crookes  tube, 
i.e.  they  are  electrons  moving  with  varying  velocities,  the 
velocity  sometimes  approaching  that  of  light.  As  the  charge 
is  negative,  they  are  deflected  by  a  magnetic  field  in  a  direction 
opposite  to  that  of  the  a-rays ;  and  as  the  mass  is  very  much 
smaller  in  proportion  to  the  charge,  the  deflection  is  much 


474  A  TEXTBOOK  OF  CHEMISTRY 

greater.  The  y-rays  seem  to  be  like  the  X-rays  or  Rontgen 
rays  and  are  probably  of  the  nature  of  ether  waves.  They  will 
penetrate  a  very  much  greater  thickness  of  metal  than  the 
/?-rays  and  these  in  turn  are  much  more  penetrating  than  the 
a-rays.  Some  of  the  radioactive  elements  give  out  all  three 
kinds  of  rays,  others  give  only  one  kind  and  still  others  two  kinds. 
The  S-rays  are  electrons  moving  much  more  slowly  than  those 
which  form  the  /?-rays.  They  may  be  given  out  from  the  sur- 
face of  radioactive  material,  or  may  be  emitted  from  any  sub- 
stance traversed  by  a-rays,  whether  solid  or  gaseous.  The 
nature  of  the  rays  furnishes  one  of  the  most  important  means  of 
identifying  different  elements. 

The  Life  of  an  Element.  In  accordance  with  the  disintegra- 
tion theory  some  of  the  atoms  of  radium  are  constantly  decom- 
posing into  helium  and  another  substance  which  was  at  first 
called  radium  emanation,  but  which  has  been  characterized  by 
Sir  William  Ramsay  as  an  element  of  the  argon  family,  the  gas 
niton.  This  has  an  atomic  weight  of  222.4,  the  difference  in 
weight  between  an  atom  of  radium  and  an  atom  of  niton  being  the 
weight  of  an  atom  of  helium.  The  rate  of  the  decomposition 
of  radium  has  been  determined  by  measuring  the  amount  of 
helium  given  in  a  number  of  weeks  or  months  by  a  given  weight 
of  radium.  The  result,  of  the  measurement  was  that  one  half 
of  a  given  quantity  of  radium  would  disintegrate  in  1760  years. 
The  rate  of  disintegration  of  niton,  on  the  other  hand,  is  so  rapid 
that  one  half  of  a  given  quantity  of  the  element  will  disintegrate 
in  3.8  days.  The  rate  of  decay  is  measured  in  this  case  by 
measuring  the  rate  at  which  the  radioactive  effect  on  the  elec- 
trometer decreases.  The  disintegration  of  uranium  to  form  the 
first  of  a  series  of  3  elements  which  are  supposed  to  stand  between 
uranium  (at.  wt.  238.5)  and  radium  (at.  wt.  226.4)  takes  place 
so  slowlv  that  one  half  of  a  given  quantity  of  the  element  would 
decompose  in  6,000,000,000  years.  These  periods,  which  are 
called  the  "  half-life  periods  "  of  the  elements,  form  one  of  the 
most  important  characteristics  of  the  radioactive  elements.  It 
is  evident  from  the  periods  which  have  been  given  for  uranium, 


RADIOCHEMISTRY  475 

radium,  and  niton  that  the  amount  of  radium  which  can  exist 
at  a  given  time  must  be  very  small  in  comparison  with  the  quan- 
tity of  uranium  in  the  world,  and  that  the  amount  of  niton  must 
always  be  very  small  in  proportion  to  the  amount  of  the  radium 
from  which  it  is  generated.  For  this  reason  Sir  William  Ramsay 
was  compelled  to  establish  the  density  and  other  properties  of 
niton  by  working  with  a  few  cubic  millimeters  of  the  gas.  He 
weighed  the  gas  with  a  microbalance,  designed  for  the  purpose, 
with  which  it  was  possible  to  weigh  innAnro  °f  a  milligram. 

Other  Radioactive  Elements.  By  the  study  of  the  rate  of  dis- 
integration of  elements  formed  by  the  decomposition  of  others  a 
considerable  number  of  radioactive  elements  have  been  identified. 
They  form  three  well-defined  series.  The  uranium  series,  begin- 
ning with  uranium  (at.  wt.  238.5)  and  probably  closing  with 
lead  (at.  wt.  207.5).  There  are  twelve  elements  between,  the 
best  characterized  being  ionium,  radium  and  niton.  The 
thorium  series  commences  with  thorium  (at.  wt.  232.4)  and  prob- 
ably closes  with  bismuth  (at.  wt.  208),  with  ten  elements  between. 
The  actinium  series  begins  with  actinium,  an  element  of  unknown 
atomic  weight,  and  closes  with  an  unknown  inactive  element, 
with  nine  elements  between.  The  half -life  of  these  elements 
varies  from  six  billions  of  years  for  uranium  and  six  hundred 
millions  of  years  for  thorium  to  a  few  seconds  for  some  of  the 
elements  derived  from  thorium  and  actinium. 

Chemical  Action  of  the  Rays.  If  radium  or  niton  is  left  in 
contact  with  water,  the  rays  which  they  emit  cause  dissociation 
of  some  of  the  water  into  oxygen  and  hydrogen.  Some  hydro- 
gen peroxide  is  also  formed.  A  glass  tube  containing  a  com- 
pound of  radium  soon  assumes  a  violet  or  brown  color.  The 
tremendous  velocity  with  which  a-particles  are  expelled  from 
radium  or  niton  gives  a  unique  and  powerful  form  of  energy, 
and  it  has  even  been  thought  that  atoms  of  other  elements  may 
be  broken  into  pieces  by  this  means  (Ramsay). 

Radiochemistry  in  Relation  to  Geology  and  Medicine.  Radio- 
active elements  are  very  widely  diffused  in  the  rocks  of  the  earth ; 
and  while  the  proportion  of  such  elements  is  very  small,  it  has 


476  A  TEXTBOOK  OF  CHEMISTRY 

been  shown  that  the  total  amount  present  in  the  crust  of  the 
earth  is  sufficient  to  account  for  the  increasing  temperature 
which  is  observed  in  deep  wells  and  mines  and  in  tunnels. 
Indeed  it  would  seem  that  the  proportion  of  radioactive  elements 
must  be  smaller  at  very  great  depths  than  it  is  near  the  surface. 
This  discovery  has  thrown  very  grave  doubts  on  estimates 
formerly  made  of  the  life  of  the  earth,  which  were  based  on  the 
supposition  that  the  earth  has  cooled  down  from  a  molten  condi- 
tion. 

The  rays  emitted  from  radium  and  other  radioactive  elements 
are  fatal  to  bacteria.  They  also  may  produce  severe  burns 
somewhat  resembling  sunburn.  They  have  been  used  with 
some  success  in  the  treatment  of  cancer,  ulcers,  lupus,  etc. 
Some  mineral  waters  in  which  chemical  analysis  has  formerly 
shown  no  peculiar  curative  substances  have  been  found  to  be 
radioactive,  and  it  seems  possible  that  beneficial  results  may  be 
obtained  by  the  use  of  such  waters. 

EXERCISES 

1.  How  many  grams  of  water  will  be  required,  theoretically,  to 
convert  a  pound  (453  g.)  of  plaster  of  Paris  into  gypsum  ? 

2.  How  much  lime  can  be  obtained  from  a  kilogram  of  marble  ? 

3.  The  heat  of  formation  : 

CO2->CaCO3 


is  42,520  small  calories  per  gram  molecule.  How  much  coal  having  the 
same  heat  of  combustion  as  carbon  would  be  required,  theoretically, 
to  prepare  one  kilogram  of  lime  ?  See  p.  27  for  the  heat  of  combustion 
of  carbon. 

4.   The  heat  of  formation  : 

BaO  +  CO2  ->  BaCO3 

is  62,220  small  calories  per  gram  molecule.  How  much  coal  having  the 
same  heat  of  combustion  as  carbon  will  be  required  to  give  one  kilogram 
of  barium  oxide  on  the  basis  of  the  reaction  : 

BaCO3  +  C  =  BaO  +  2  CO  ? 


ALKALI-EARTH  METALS  477 

5.  One  hundred  cubic  centimeters  of  a  natural  water  require  6  cc. 
of  tenth  normal  sodium  hydroxide  to  give  a  pink  color  with  phenol- 
phthalein  and  the  solution  then  requires  10  cc.  of  tenth  normal  hydro- 
chloric acid  to  give  a  red  color  with  methyl  red.     How  much  carbonic 
acid  and  how  much  bicarbonate,  calculated  as  calcium  bicarbonate,  are 
present  in  one  liter  of  the  water  ? 

6.  How  much  lime  (CaO)  will  be  required  to  furnish  enough  calcium 
hydroxide  to  soften  one  U.  S.  gallon  (3.785  liters)  of  the  water  just 
referred  to  ? 


CHAPTER  XXVII 

ALTERNATE    METALS  OF   GROUP   II.      MAGNESIUM,   ZINC, 
CADMIUM    AND    MERCURY 

IN  its  occurrence  and  in  the  properties  of  its  most  common 
compounds,  magnesium  resembles  calcium  rather  than  zinc,  but 
in  the  metallic  form  the  resemblance  to  zinc  is  more  marked. 
Magnesium  decomposes  water  slowly  at  the  boiling  point,  while 
zinc  and  cadmium  decompose  it  readily  at  a  higher  temperature. 
Mercury  in  many  of  its  properties  seems  to  be  more  closely 
related  to  copper  and  silver  than  to  cadmium,  and  some  authors 
formerly  placed  it  in  the  first  group.  It  does  not  decompose 
water  at  any  temperature  and  it  forms  compounds  in  which  it  is 
apparently  univalent  as  well  as  those  in  which  it  is  bivalent. 
Magnesium,  zinc  and  cadmium  are  always  bivalent. 

Magnesium,  Mg,  24.32,  is  found  as  the  carbonate,  magnesite, 
MgCOs,  as  a  double  carbonate  of  calcium  and  magnesium, 
dolomite,  CaCOs.MgCOs,  as  the  double  chloride  with  potassium, 
carnallite,  KCl.MgCl2.6  H2O,  and  as  a  principal  constituent  of 
many  silicates,  especially  talc  or  soapstone,  MgsH^SiOs)^ 
serpentine,  Mg3Si2O7.2  H2O,  and  meerschaum,  Mg2Si3O8.2  H2O. 
The  sulfate,  Epsom  salts,  MgSO4.7  H2O,  and  the  chloride, 
MgCl2.6  H2O,  are  also  found  in  many  natural  waters. 

Preparation,  Properties.  Metallic  magnesium  is  obtained 
by  the  electrolysis  of  fused  carnallite,  MgCl2.KCl,  from  which 
the  water  of  hydration  has  been  expelled  by  heat.  It  is  a 
silver-white,  very  light  metal,  having  a  specific  gravity  of  only 
1.75,  slightly  lower  than  that  of  beryllium.  It  melts  at  651° 
and  boils  at  about  1100°.  Magnesium  wire  or  ribbon  burns  in 
the  air  with  a  very  intense  white  light  that  is  particularly  rich 
in  the  more  refrangible  rays,  which  affect  the  photographic  plate. 

478 


GROUP  II:    MAGNESIUM  479 

For  this  reason  powdered  magnesium  is  the  effective  constituent 
of  flash-light  powders.  It  is  estimated  that  10  per  cent  of  the 
energy  of  burning  magnesium  appears  as  light,  a  very  much 
larger  per  cent  than  is  secured  by  any  ordinary  illuminant. 
The  temperature  of  burning  magnesium  is  not,  'however,  very 
high  —  only  about  1340°.  Magnesium  tarnishes  only  slightly, 
if  at  all,  in  dry  air  and  very  slowly  in  moist  air,  so  that  it  can 
be  kept  indefinitely  without  special  precautions. 

Magnesium  is  used  in  the  laboratory  as  a  powerful  reducing 
agent  for  the  preparation  of  silicon  and  boron.  It  is  also  used 
in  a  great  variety  of  syntheses  of  organic  compounds. 

Magnesium  Oxide,  MgO,  is  most  easily  prepared  by  heating 
magnesium  carbonate,  MgCO3,  which  decomposes  at  a  much 
lower  temperature  than  calcium  carbonate.  It  is  a  light,  white 
powder  often  called  magnesia  usta,  or  burnt  magnesia.  It  is 
infusible  in  any  ordinary  furnace,  but  may  be  volatilized  in  the 
electric  furnace.  It  is  used  for  crucibles  and  for  some  forms  of 
apparatus  which  must  withstand  extremely  high  temperatures. 
It  is  used  as  a  basic  lining  for  metallurgical  furnaces,  especially 
for  the  basic  process  for  steel  (p.  548). 

Magnesium  Hydroxide,  Mg(OH)2,  is  obtained  as  a  white 
precipitate  on  the  addition  of  sodium  hydroxide,  NaOH,  or  barium 
hydroxide,  Ba(OH)2,  to  a  solution  containing  almost  any  soluble 
salt  of  magnesium.  It  is  much  less  soluble  than  calcium  hydrox- 
ide, dissolving  in  about  6500  parts  of  water.  In  spite  of  this 
difficult  solubility,  however,  it  is  not  precipitated  in  the  presence 
of  ammonium  salts.  This  is  due  probably  to  two  reasons : 
(1)  in  the  presence  of  an  ammonium  salt  the  ionization, 

NH4OH  ^±  NH4+  +  OH- 

is  repressed  by  the  ammonium  ions,  NH4+,  of  the  salt,  and  the 
concentration  of  the  hydroxide  ions,  OH",  is  low  in  such  a  solu- 
tion; and  (2)  because  magnesium  hydroxide  as  a  unibivalent  com- 
pound very  probably  forms  intermediate  ions,  MgOH+  +  OH", 
which  interfere  with  the  application  of  the  ordinary  law  of  the 
solubility  product  (p.  377). 


480  A  TEXTBOOK  OF  CHEMISTRY 

Magnesium  hydroxide  is  easily  decomposed  by  heat  into 
magnesium  oxide  and  water. 

Magnesium  Chloride,  MgCl2.6H2O,  crystallizes  from  a  con- 
centrated solution  of  the  salt.  It  is  very  easily  soluble  in 
water.  When  "an  attempt  is  made  to  drive  out  the  water  of  the 
salt  by  heating  it,  both  water  and  hydrochloric  acid  escape,  and 
a  mixture  of  variable  composition,  containing  chiefly  mag- 
nesium oxide,  finally  remains : 

MgCl2  +  H2O  =  MgO  +  2  HC1 

The  process  has  been  used  to  a  limited  extent  as  a  basis  for  the 
preparation  of  hydrochloric  acid  and  chlorine. 

Magnesium  Ammonium  Chloride,  MgNH4Cl3.6  H2O.  This 
salt  is  easily  prepared  by  crystallizing  from  water  a  mixture  of 
equimolecular  amounts  of  magnesium  chloride  and  ammonium 
chloride.  The  water  of  hydration  may  be  expelled  with  very 
little  loss  of  hydrochloric  acid,  and  on  heating  the  anhydrous 
salt  to  a  slightly  higher  temperature  the  ammonium  chloride 
dissociates  and  escapes,  leaving  anhydrous  magnesium  chloride 
behind. 

Magnesium  Sulfate,  MgSO4.7  H2O,  or  Epsom  Salts,  is  found 
in  some  mineral  waters  used  for  their  medicinal  properties,  es- 
pecially in  Hunyadi  water,  in  which  it  is  associated  with  sodium 
sulfate,  Na2SO4. 

*  Magnesium  Sulfide,  MgS,  may  be  prepared  by  heating  a 
mixture  of  magnesium  and  sulfur.     It  is  decomposed  by  water, 
giving  magnesium  hydroxide  and  hydrogen  sulfide : 

MgS  +  2  HOH  =  Mg(OH)2  +  H2S 

Both  the  insolubility  of  the  hydroxide  and  the  volatility  of  the 
hydrogen  sulfide  contribute  to  cause  the  reaction  to  go  to  com- 
pletion. 

*  Magnesium  Ammonium  Phosphate,  MgNH4PO4.6  H2O,  is 
a  difficultly  soluble  salt  which  is  formed  when  solutions  contain- 
ing magnesium,  ammonium  and  a  soluble  phosphate  are  brought 
together.     It  is  used  in  analytical  chemistry  for  the  determina- 
tion of  both  magnesium  and  phosphoric  acid.     A  precipitate 


GROUP  II:    ZINC  481 

having  the  exact  composition  represented  by  the  formula  can 
be  obtained  only  by  securing  exactly  the  right  conditions  as 
regards  the  concentration  of  the  various  solutions  employed. 
When  heated,  the  compound  decomposes  quantitatively  with  the 
formation  of  magnesium  pyrophosphate,  Mg2P2O7. 

Zinc,  Zn,  65.37.  Occurrence.  Zinc  is  found  in  nature  as  the 
sulfide,  sphalerite,  ZnS,  the  carbonate,  smithsonite,  ZnCOs,  the 
silicates,  willemite,  Zn2SiO4,  and  calamine,  H^ZnSiO-i,  and  in  the 
mineral  franklinite,  which  consists  of  oxides  of  iron,  zinc  and 
manganese.  Franklinite  is  an  important  ore  in  New  Jersey  and 
is  used  as  a  source  of  zinc,  manganese  and  iron. 

Metallurgy.  The  sulfide  of  zinc  is  converted  into  the  oxide 
by  roasting  it,  that  is  by  heating  it  in  a  furnace  with  free  access 
of  air : 

ZnS  +  3  O  =  ZnO  +  SO2 

The  oxide,  obtained  in  this  way  or  by  heating  the  carbonate, 
is  mixed  with  coal  and  heated  to  a  high  temperature  in  an  earth- 
enware retort  having  a  receiver  luted  to  it  with  clay.  The  zinc 
oxide  is  reduced,  and  the  zinc,  which  boils  at  925°,  distils  over. 
Zinc  melts  at  419.4°. 

Impure  zinc  dissolves  easily  in  hydrochloric  or  in  dilute  sul- 
furic  acid,  with  evolution  of  hydrogen.  Pure  zinc,  which  can 
be  obtained  by  distillation  in  a  vacuum,  is  attacked  very  slowly 
or  not  at  all  by  these  acids,  but  dissolves  readily  in  contact  with 
platinum.  It  has  been  pointed  out  that  these  facts  indicate  that 
the  solution  is  always  associated  with  electrical  phenomena. 
Impure  zinc  which  is  covered  by  a  thin  film  of  amalgam,  giving 
it  a  homogeneous  surface,  is  also  not  attacked  by  the  dilute  acids, 
and  such  amalgamated  zinc  is  used  in  electric  batteries.  Zinc 
has  a  specific  gravity  of  6.9. 

Uses.  Galvanized  Iron.  Metallic  zinc  is  used  chiefly  in  brass, 
an  alloy  of  the  metal  with  about  twice  its  weight  of  copper,  and 
as  a  coating  for  iron  to  protect  it  from  rusting.  It  is  also  a  con- 
stituent of  many  of  the  bronzes,  and  is  used  as  the  metal  which  is 
dissolved  or  corroded  in  most  forms  of  primary  electrical  batter- 


482  A  TEXTBOOK  OF  CHEMISTRY 

ies,  especially  in  the  gravity  cell  used  in  telegraphy  and  in  the 
so-called  "  dry  "  batteries. 

"  Galvanized  "  iron  is  prepared  by  dipping  carefully  cleaned 
sheet  iron  or  other  iron  or  steel  articles  in  melted  zinc.  The 
value  of  the  coating  depends  on  two  properties :  first,  zinc  is 
electropositive  with  reference  to  iron  and  when  the  two  metals 
are  in  contact  with  each  other  and  also  in  contact  with  an  elec- 
trolyte the  zinc  is  attacked  and  the  iron  is  protected ;  second,  the 
action  upon  the  zinc  causes  the  formation  of  a  very  thin, 
coherent  coating  of  zinc  oxide  or  hydroxide,  which  is  practically 
insoluble  in  water  and  protects  the  zinc  from  further  action. 
A  small  amount  of  zinc  passes  into  solution,  however,  and  this 
may  be  increased  very  considerably  in  the  presence  of  even 
weak  acids.  As  zinc  salts  are  poisonous,  pails  or  dishes  of  gal- 
vanized iron  are  not  suitable  for  culinary  use. 

The  coating  of  zinc  has  been  sometimes  applied  to  the  iron  by 
an  electrolytic  method,  and  the  term  "  galvanized  iron  "  came 
from  this  method  of  manufacture. 

Sherardized  Iron.  A  new  process  for  coating  iron  with  zinc 
has  been  developed  by  Sherard  Cowper-Cowles.  The  articles  to 
be  coated  are  heated  with  zinc  dust  in  iron  drums  at  500°-600° 
for  thirty  minutes  to  several  hours,  according  to  the  thickness  of 
the  coating  desired.  The  process  is  somewhat  analogous  to  the 
manufacture  of  cementation  steel.  See  Johnson  and  Woolrich, 
Trans.  Am.  Electrochem.  Soc.  21,  561  (1912).  Cowper-Cowles, 
Electrochem.  and  Met.  Ind.  6,  189  (1908). 

Zinc  Oxide,  ZnO,  is  prepared  by  burning  the  vapors  of  metallic 
zinc.  It  is  a  white  powder  and  gives  with  linseed  oil  an  excellent 
pigment,  which  has  the  advantage  of  not  being  blackened  by 
hydrogen  sulfide  because  zinc  sulfide  is  also  white.  Zinc  oxide 
is  often  used  with  phosphoric  acid  for  a  cement  in  dental  work. 
The  two  substances  combine  to  form  a  basic  zinc  phosphate 
which  sets  to  a  hard  mass  that  adheres  strongly  to  the  surfaces 
with  which  it  is  in  contact.  Zinc  oxide  is  yellow  when  hot,  but 
turns  white  again  on  cooling,  a  property  used  for  the  detection 
of  zinc  compounds  in  blowpipe  reactions. 


GROUP  II:    ZINC,  CADMIUM  483 

Zinc  Hydroxide,  Zn(OH)2,  forms  as  a  white  precipitate  on 
adding  a  soluble  hydroxide  to  a  solution  of  a  zinc  salt.  It  dis- 
solves in  an  excess  of  sodium  hydroxide  or  potassium  hydroxide, 
forming  sodium  zincate,  Na2ZnC>2,  or  potassium  zincate,  K^ZnC^. 
In  forming  these  compounds  zinc  hydroxide  seems  to  act  as  an 
acid,  which  might  be  called  zincic  acid,  and  the  formula  might  be 
written  H2ZnO2.  Toward  acids,  however,  zinc  hydroxide  con- 
ducts itself  as  a  true  hydroxide  or  base.  Compounds  which 
exhibit  a  dual  nature  of  this  sort,  acting  in  some  conditions  as 
acids  and  in  others  as  bases,  are  said  to  be  amphoteric. 

Zinc  Chloride,  ZnCl2.  An  aqueous  solution  of  this  salt  is 
easily  prepared  by  dissolving  metallic  zinc  or  zinc  oxide  in  hydro- 
chloric acid.  Unlike  magnesium  chloride,  the  solution  loses 
only  a  small  amount  of  hydrochloric  acid  when  heated  to  a  high 
temperature  to  expel  the  water.  The  pure,  anhydrous  salt 
melts  at  290°-297°  and  boils  at  730°.  When  boiled  in  an  iron 
tube,  it  furnishes  an  easy  means  of  securing  a  constant,  rather 
high  temperature  and  has  been  found  useful  for  this  purpose. 
Zinc  chloride  is  used  in  the  treatment  of  wooden  ties  to  prevent 
decay. 

Zinc  Sulfate,  ZnSO4.7  H2O,  or  White  Vitriol.  Anhydrous 
zinc  sulfate,  ZnSO4,  can  be  prepared  by  roasting  the  sulfide,  ZnS, 
at  a  moderate  temperature.  At  a  higher  temperature  the  sul- 
fide roasts  to  the  oxide  and  sulfur  dioxide.  The  hydrate, 
ZnSO4.7  H2O,  forms  rhombic  crystals  and  is  easily  soluble. 

Zinc  Sulfide,  ZnS,  forms  as  a  white  precipitate  when  hydrogen 
sulfide  is  passed  into  a  neutral  or  alkaline  solution  of  a  zinc  salt. 
The  precipitate  is  formed  even  in  slightly  acid  solutions,  and  care- 
ful attention  must  be  paid  to  the  amount  and  character  of  the 
acid  present  if  a  separation  from  other  metals  is  desired.  In 
the  presence  of  sulfuric  acid  which  is  weaker  than  fifth  normal 
(about  1  per  cent)  the  sulfide  will  be  precipitated. 

Cadmium  (Cd,  112.40).  Many  zinc  ores  contain  a  small 
amount  of  cadmium.  As  the  boiling  point  of  cadmium  (785°) 
is  considerably  lower  than  that  of  zinc  (925°),  the  former  distills 
over  first  in  the  preparation  of  zinc,  and  by  collecting  these  por- 


484  A  TEXTBOOK  OF  CHEMISTRY 

tions  and  subjecting  them  to  fractional  distillation,  nearly  pure 
cadmium  can  be  prepared. 

Metallic  cadmium  closely  resembles  zinc  in  appearance  and 
in  many  of  its  properties.  It  melts  at  320.9°,  boils  at  785°  and  has 
a  specific  gravity  of  8.65.  It  was  formerly  used  in  amalgams 
for  filling  teeth,  but  other  amalgams  are  now  considered  more 
suitable.  It  is  a  constituent  of  Wood's  metal  and  of  the  easily 
fusible  alloys  used  for  safety  fuses  in  electrical  circuits  and  for 
automatic  sprinklers  used  for  protection  against  fire. 

*  Cadmium  Hydroxide,  Cd(OH)2,  is  easily  obtained  as  a  white 
precipitate.    It  dissolves  easily  in  acids  but  does  not  dissolve  in 
solutions  of  sodium  or  potassium  hydroxides,  as  zinc  hydroxide 
does.     It  is  decomposed  when  heated,  giving  cadmium  oxide, 
CdO,  as  a  brown  powder. 

*  Cadmium  Sulfate,  3  CdSO4.8  H2O,  is  an  easily  soluble  salt 
used  in  the  Weston  standard  cells,  which  are  the  most  satisfac- 
tory primary  standard  for  the  measurement  of  electromotive 
force. 

Cadmium  Sulfide,  CdS,  forms  as  a  yellow  precipitate  in  solu- 
tions of  cadmium  salts  which  do  not  contain  too  much  free  acid 
or  too  much  of  salts  which  interfere  with  the  precipitation. 
From  a  solution  containing  both  zinc  and  cadmium  in  which 
sulf uric  acid  is  present  and  the  concentration  of  the  hydrogen  ion 
is  between  fifth  normal  and  twice  normal  and  other  interfering 
salts  or  acids  are  absent,  the  precipitation  of  the  cadmium  is 
practically  complete,  while  only  a  small  amount  of  the  zinc  will 
come  down.  For  a  complete  separation,  however,  the  cadmium 
sulfide  must  be  dissolved  and  reprecipitated.  In  hydrochloric 
acid  stronger  than  0.3  normal  cadmium  sulfide  is  not  completely 
precipitated.  Cadmium  sulfide  dissolves  readily  in  boiling, 
dilute  sulfuric  acid,  but  is  insoluble  in  a  solution  of  potassium 
cyanide,  KCN. 

Mercury,  Hg,  200.6.  Occurrence.  Metallurgy.  From  the 
positions  of  the  elements  of  Group  II  in  the  electromotive 
series  (p.  436)  mercury  is  the  only  element  of  the  group  which 
could  appear  in  the  free  state  in  nature.  It  is  occasionally  found 


GROUP  II:    MERCURY  485 

in  small  globules  disseminated  in  porous  rocks.  Mercury  occurs 
chiefly,  however,  in  the  form  of  the  native  sulfide,  cinnabar,  HgS, 
a  brilliant  red  mineral,  when  pure.  When  ores  containing  cinna- 
bar are  roasted  by  heating  in  a  current  of  air,  the  sulfur  burns  to 
sulfur  dioxide  while  the  mercury  distills  and  is  condensed  in 
long  flues  where  the  vapors  must  be  very  thoroughly  cooled  to 
prevent  loss:  HgS  +  O2  =  Hg  +  SO2 

Mercury  may  also  be  obtained  by  mixing  the  sulfide  with  lime 
or  with  iron  and  distilling : 

2  CaO  +  2  HgS  =  2  CaS  +  2Hg+  O2 
Fe  +  HgS  =  FeS  +  Hg 

Mercury  can  be  purified  by  allowing  it  to  fall  in  very  minute 
globules,  through  a  chamois  skin  tied  over  the  end  of  a  glass 
funnel,  into  dilute  nitric  acid  contained  in  a  tube  2  meters  long. 
The  tube  is  drawn  out  and  bent  upward  at  the  bottom  so  that  a 
short  column  of  mercury  in  the  overflow  tube  balances  the  col- 
umn of  nitric  acid.  The  nitric  acid  dissolves  zinc,  arsenic,  lead 
and  nearly  all  of  the  other  metals  likely  to  be  present. 

Mercury  may  also  be  separated  from  nearly  all  other  metals 
by  distillation  under  diminished  pressure.  If  a  very  little  air 
is  allowed  to  pass  through  the  mercury  by  means  of  a  very  fine, 
hairlike,  capillary  tube,  troublesome  bumping  of  the  mercury 
can  be  avoided  and  zinc  and  some  other  metals  are  oxidized, 
giving  purer  mercury  than  if  the  distillation  is  carried  out  in 
the  absence  of  air  (Hulett). 

Properties  and  Uses.     Mercury  is  a  heavy,  mobile  liquid,  with 

O°  9O° 

a  density,  at  ^  =  13.5956  or  at  ^-  =  13.5463.     It  freezes  at 

—  38.70°.     Its  freezing  point  on  a  Fahrenheit  thermometer  is 

—  37.7°.     It  will  be  noticed  that  the  scales  of  the  Centigrade 
and  Fahrenheit  thermometers  approach  very  closely  together 
at  this  temperature.     Mercury  boils  at  357°.     Its  critical  tem- 
perature is  about  1275°,  and  its  critical  pressure  is  calculated  as 
about  675  atmospheres,  an  extraordinarily  high  value.     (Menzies, 


486  A  TEXTBOOK  OF  CHEMISTRY 

J.  Am.  Chem.  Soc.,  Sept.,  1913.  Konigsberger,  Chem.  Ztg.  13d, 
1321  (1913) .)  Mercury  oxidizes  slowly  to  red  mercuric  oxide,  HgO, 
when  heated  to  its  boiling  point  in  the  air.  (See  Lavoisier's 
experiment,  p.  19.)  Mercury  dissolves  very  slowly  to  mercur- 
ous  nitrate,  HgNOs,  in  dilute  nitric  acid,  with  evolution  of  nitric 
oxide,  NO.  It  is  insoluble  in  hydrochloric  acid,  but  is  converted 
into  mercurous  sulfate,  Hg2SO4,  by  hot,  concentrated  sulfuric 
acid,  with  evolution  of  sulfur  dioxide. 

Mercury  is  used  in  the  amalgamation  processes  for  the  re- 
covery of  gold  and  silver  (p.  441),  in  making  thermometers,  for 
barometers  and  manometers,  in  mercury  air  pumps  and  in  the 
collection  and  measurement  of  gases.  Its  advantages  over  all 
other  substances  used  in  thermometers  are,  especially,  that  it 
does  not  wet  or  attack  the  glass,  that  it  is  liquid  over  a  wide 
range  of  temperature,  including  the  common  range  of  air  tem- 
peratures, and  that  its  rate  of  expansion  is  very  uniform.  Its 
coefficient  of  expansion  between  0°  and  100°  is  so  nearly  constant 
that  a  mercury  thermometer  graduated  in  equal  degrees  does 
not  differ  from  the  standard  hydrogen  scale  by  more  than  0.2° 
at  any  point  between  these  temperatures.  Mercury  thermom- 
eters cannot,  of  course,  be  used  at  temperatures  below  —  39° 
and  ordinary  thermometers  cannot  be  used  above  300°  — 
indeed,  the  thread  of  an  ordinary  thermometer  will  usually  break 
before  that  temperature  is  reached.  By  filling  the  space 
above  the  mercury  with  nitrogen  under  pressure,  however, 
thermometers  graduated  to  460°  are  made,  and  by  filling  the 
space  with  carbon  dioxide  the  range  has  been  carried  to  550°  or 
above.  There  is  likely  to  be  a  large  zero-point  correction  for 
such  thermometers,  and  they  must  be  carefully  treated,  if 
accurate  results  are  required.  The  stem  correction  is  also 
large,  unless  the  whole  thermometer  is  immersed  in  the  sub- 
stance whose  temperature  is  to  be  measured. 

Amalgams.  The  alloys  of  mercury  are  called  amalgams. 
Many  metals,  such  as  sodium,  potassium,  copper,  silver,  gold, 
zinc,  cadmium,  tin  and  lead,  dissolve  in  or  alloy  with  mercury 
in  all  proportions  or  give  amalgams  having  a  wide  range  in  their 


GROUP  II:  AMALGAMS 


487 


composition.  Other  metals,  as  iron  and  platinum,  dissolve  in 
mercury  to  only  a  trifling  extent  or  not  at  all.  The  amalgams  of 
gold  and  silver  are  used  to  separate  these  metals  from  large 
masses  of  other  substances  mixed  with  them  in  their  ores. 
Sodium  amalgam  is  often  used  as  a  reducing  agent,  especially 
for  organic  compounds.  Zinc  amalgam  and  other  amalgams 
may  be  used  in  the  same  way.  An  amalgam  of  tin  was  formerly 
used  for  the  backs  of  mirrors,  but  has  been  replaced  by  a  thin 
film  of  metallic  silver  in  modern  mirrors.  An  amalgam  with 
silver  and  other  metals  is  used  for  filling  teeth. 

In  many  cases  mercury  combines  with  metals  to  form  definite 
compounds.     Such  compounds  are  most  easily  identified  by  a 


sou 
300° 
250° 

82°° 
H 

H  150° 

3 

|  100° 

H 
P 

50° 
0° 

-50 

Per  Cent.) 
Mercury  ) 
Per  Cent.  ) 

II 

\ 

' 

/ 

\ 

7 

\\ 

\\ 

-^^. 

^-^ 

---^. 

—^^^^ 

/ 

1 

^ 

-—  -^ 

X 

1 

1 

0          10         20         30         40         50         60         70         80         90       100 
LOO      90         80         70         60         50         40         30         2O         10          0 

Fig.  102 

study  of  the  freezing  point  curve  of  amalgams  of  varying  composi- 
tion. In  Fig.  102  the  ordinates  give  the  melting  points  and  the 
abscissas  give  the  composition  of  a  series  of  amalgams  of  mercury 
with  sodium.  It  will  be  seen  from  the  figure  that  the  addition 


488  A  TEXTBOOK  OF  CHEMISTRY 

of  mercury  to  sodium  lowers  its  melting  point  until  a  minimum 
is  reached  for  an  amalgam  containing  about  40  per  cent  of  mer- 
cury and  60  per  cent  of  sodium.  This  minimum  is  called  a 
eutectic  point.  An  amalgam  of  this  composition  melts  at  21°. 
Further  addition  of  mercury  raises  the  melting  point  till  a  maxi- 
mum is  reached  for  an  amalgam  containing  5.4  percent  of  sodium 
and  94.6  per  cent  of  mercury,  which  melts  at  346°.  Further 
addition  of  mercury  lowers  the  melting  point  till  this  would, 
undoubtedly,  fall  below  the  melting  point  of  pure  mercury.  A 
compound  having  the  formula  NaHg2  would  contain  5.43  per 
cent  of  sodium  and  94.57  per  cent  of  mercury.  Evidently  the 
amalgam  of  the  highest  melting  point  is  a  compound  of  this 
formula.  The  addition  of  either  sodium  or  mercury  to  this 
compound  lowers  its  melting  point  just  as  the  addition  of  salt  or 
any  soluble  substance  lowers  the  melting  point  of  ice.  Changes 
in  the  direction  of  the  curve  at  other  points  indicate  that  other 
compounds  of  mercury  and  sodium  are  present  in  some  of  the 
amalgams,  but  these  details  are  not  shown  in  the  figure.  See 
Kurnakow,  Z.  anorg.  Chem.  23,  443  (1900). 

Compounds  of  Mercury.  Mercury  forms  mercurous  com- 
pounds, such  as  Hg2O,  Hg2Cl2,  Hg2SO4,  in  which  it  appears  uni- 
valent,  but  in  which  it  is  probably  really  bivalent  as  expressed 

Hg-Cl 

by  the  graphical  formula,    |  .    It  also  forms  mercuric  com- 

Hg-Cl 

pounds,  such  as  HgO,  HgCl2,  HgSO^  in  which  it  is  clearly  bi- 
valent. In  the  formation  of  these  two  classes  of  compounds  and 
also  in  its  conduct  toward  nitric,  sulfuric  and  hydrochloric  acids 
mercury  resembles  copper  rather  than  zinc  or  cadmium,  and  it 
has  sometimes  been  classified  under  the  first  group  of  the  Periodic 
System  in  place  of  gold. 

Mercurous  Oxide,  Hg2O,  is  formed  as  a  black  precipitate 
when  a  solution  of  sodium  hydroxide  is  added  to  a  solution  of 
mercurous  nitrate,  HgNO3,  or  when  calomel,  Hg2Cl2,  is  digested 
with  a  solution  of  sodium  hydroxide. 

Mercuric  Oxide,  HgO,  is  formed  slowly  as  a  heavy  red  crys- 


GROUP  II:    MERCURY  489 

talline  powder  when  mercury  is  heated  to  its  boiling  point  in  the 
air.  It  is  obtained  more  easily  by  heating  the  nitrate.  A  yellow 
precipitate  having  the  same  composition  is  formed  on  adding  an 
alkali  to  a  solution  of  a  mercuric  salt. 

Mercuric  Sulfide,  HgS.  The  mineral  cinnabar,  HgS,  is  a 
bright  red  compound.  When  hydrogen  sulfide  is  passed  into  a 
solution  of  a  mercuric  salt  a  black  mercuric  sulfide  of  exactly 
the  same  composition  is  precipitated.  By  subliming  the  black 
sulfide,  or  by  warming  it  with  a  solution  of  sodium  sulfide,  it 
can  be  converted  into  the  red  variety.  The  red  form  is  used 
under  the  name  of  vermilion  as  a  brilliant  red  pigment.  When 
applied  to  iron  or  zinc,  however,  it  is  decomposed  with  libera- 
tion of  metallic  mercury. 

Mercurous  Chloride,  or  Calomel,  Hg2Cl2,  is  prepared  by  sub- 
liming a  mixture  of  mercuric  chloride,  HgCl2,  and  mercury,  or  a 
mixture  of  mercuric  sulfate,  HgSC>4,  salt  and  mercury.  The 
crude  product  usually  contains  a  little  mercuric  chloride,  which  is 
removed  by  treatment  with  alcohol,  in  which  the  mercurous 
chloride  is  insoluble  while  the  mercuric  chloride  is  easily  soluble. 
The  gram  molecular  volume  of  the  vapor  of  mercurous  chloride 
weighs  about  236  grams,  corresponding  to  the  formula  HgCl, 
but  it  has  been  shown  that  the  vapor  really  consists  of  a  mixture 
of  mercuric  chloride  and  mercury  (HgCl2  +  Hg)  (Alex.  Smith, 
J.  Am.  Chem.  Soc.  32  1541  (1910)). 

From  this  it  seems  probable  that  the  true  formula  of  mercurous 
chloride  is  Hg2Cl2. 

Calomel  is  used  as  a  medicine.  It  is  now  usually  administered 
in  very  small  doses  and  mixed  with  sodium  bicarbonate,  NaHCO3, 
to  render  it  less  soluble  in  the  acid  gastric  juice.  In  former  times 
the  careless  administration  of  large  doses  sometimes  caused  sali- 
vation and  other  serious  injuries  to  patients. 

Mercurous  chloride  is  formed  as  a  white  precipitate  on  adding 
hydrochloric  acid  or  a  soluble  chloride  to  a  solution  of  mercurous 
nitrate  or  of  some  other  soluble  mercurous  salt. 

Mercuric  Chloride  or  Corrosive  Sublimate,  HgCl2,  is  prepared 
by  subliming  a  mixture  of  mercuric  sulfate,  HgS(>4,  and  salt, 


490  A  TEXTBOOK  OF  CHEMISTRY 

NaCl.  It  is  a  white,  crystalline  salt,  which  melts  at  265°  and 
boils  at  307°.  It  is  soluble  in  about  14  parts  of  cold  water  and 
more  easily  soluble  in  alcohol.  It  also  dissolves  in  ether.  When 
taken  internally  it  is  very  poisonous.  The  best  antidote  is  the 
white  of  an  egg,  with  which  it  forms  an  insoluble  compound. 
The  solution  in  alcohol  is  sometimes  used  as  a  poison  for  insects. 
A  dilute  solution  (usually  1 :  1000)  is  much  used  as  an  antiseptic 
in  surgery.  From  such  a  solution  the  mercury  does  not  seem  to 
be  absorbed  from  a  wound  or  through  the  skin. 

Mercuric  Iodide,  Hgl2,  is  precipitated  as  a  scarlet  powder 
on  adding  a  solution  of  potassium  iodide,  KI,  to  a  solution  of 
mercuric  chloride.  The  precipitate  dissolves  in  an  excess  of  the 
potassium  iodide,  forming  the  complex  salt,  K2HgI4.  Sodium 
hydroxide  produces  no  precipitate  in  such  a  solution,  evidently 
because  it  contains  only  a  very  small  number  of  mercuric  ions, 
Hg++.  An  alkaline  solution  prepared  in  this  manner  is  used 
under  the  name  of  Nessler's  solution  as  an  extremely  sensitive 
reagent  for  ammonia. 

Mercurous  Nitrate,  HgNO3  or  Hg2(NO3)2,  is  formed  by  the 
solution  of  mercury  in  cold,  dilute  nitric  acid.  It  is  hydrolyzed 
by  water,  giving  a  basic  nitrate,  Hg2(OH)NO3,  hence  to  secure 
a  clear  solution  a  little  nitric  acid  must  be  added  to  carry  the 
reversible  reaction : 

Hg2(NO3)2  +  HOH  ^±  Hg2(OH)NO3  +  HNO3 
to  the  left.     To  counteract  the  oxidation  to  mercuric  nitrate, 
Hg(NOs)2,  by  the  oxygen  of  the  air,  some  metallic  mercury  must 
be  kept  in  contact  with  the  solution  : 

2  Hg2(NO3)2  +  4  HNO3  +  O2  =  4  Hg(NO3)2  +  2  H2O 
Hg(N03)2  +  Hg  =  Hg2(N03)2 

*  Mercuric  Nitrate,  Hg(NO3)2.8  H2O,  is  obtained  by  dissolv- 
ing mercury  in  warm,  concentrated  nitric  acid. 

*  Mercuric  Cyanide,  Hg(CN)2,  can  be  prepared  by  dissolving 
precipitated  mercuric  oxide  in  a  solution  of  hydrocyanic  acid, 
HCN.     It    decomposes    into    mercury    and    cyanogen,    C2N2, 
when  heated. 


MAGNESIUM,   ZINC,  CADMIUM  AND  MERCURY     491 

*  Mercuric  Fulminate,  Hg(ONC)2,  is  used  in  cartridges  and 
percussion  caps  for  firearms  and  in  detonating  caps  for  firing 
dynamite  and  nitroglycerin. 

It  is  hydrolyzed  by  hydrochloric  acid  and  water  to  hydroxyl- 
amine  hydrochloride,  NH2OH.HC1,  and  formic  acid,  HCO2H. 


Hg<  +4HC1  +  4H20 

^—C 


=  HgCl2  +  2  H-0— NH2.HC1  +  2  O= 


lonization  of  Compounds  of  Cadmium  and  Mercury.  For 
some  reason,  not  understood,  the  chlorides  and  sulfates  of  cad- 
mium and  mercuric  mercury  ionize  to  a  much  smaller  degree 
than  the  corresponding  salts  of  most  other  metals. 

Solubility  of  the  Sulfides  of  Group  II.  The  sulfides  of  the 
metals  of  the  first  division  of  Group  II,  CaS,  SrS  and  BaS,  are 
hydrolyzed  by  water,  forming  hydroxides  and  soluble  hydro- 

sulfides : 

2  CaS  +  2  HOH  =  Ca(OH)2  +  Ca(SH)2 

Magnesium  sulfide,  MgS,  gives  with  water  magnesium  hy- 
droxide, Mg(OH)2,  and  hydrogen  sulfide.  Zinc  sulfide,  ZnS,  is 
not  affected  by  water,  but  dissolves  in  strong  acids,  if  not  too 
dilute.  It  is  almost  insoluble  in  such  a  weak  acid  as  acetic  acid. 
Cadmium  sulfide,  CdS,  dissolves  readily,  especially  on  warming,  in 
moderately  concentrated,  strong  acids,  especially  in  nitric  acid 
(5  per  cent)  or  sulfuric  acid  (15  per  cent).  Mercuric  sulfide,  HgS 
does  not  dissolve,  even  in  boiling  nitric  acid,  but  dissolves 
easily  in  aqua  regia.  These  relations  furnish  a  ready  means  of 
separating  magnesium,  zinc,  cadmium  and  mercury  from  each 
other  and  cause  them  to  be  classified  in  three  different  groups 
for  analytical  purposes. 

Conduct  of  Solutions  of  Magnesium,  Zinc  and  Cadmium 
Salts  toward  Ammonium  Hydroxide.  Ammonium  hydroxide 
gives  no  precipitate  with  salts  of  these  metals  in  solutions  con- 


492  A  TEXT  BOOK  OF   CHEMISTRY 

taining  ammonium  chloride.  The  zinc  and  cadmium  salts  form 
complex  compounds,  which  are  soluble,  such  as  Zn(NHs)4SO4 
and  Cd(NHs)4Cl2.  These  resemble  the  corresponding  com- 
pounds of  copper  but  are  colorless.  For  Mg  see  p.  479. 

Ammono-mercuric  Compounds.  When  a  solution  of  an  alka- 
line hydroxide  is  added  to  a  solution  of  a  mercurous  or  mercuric 
salt,  mercurous  oxide,  Hg2O,  or  mercuric  oxide,  HgO,  is  precipi- 
tated, as  has  been  stated.  If  ammonium  hydroxide,  NH^OH,  is 
added  to  such  a  solution,  however,  compounds  of  a  wholly  differ- 
ent type,  called  ammonobasic  mercuric  compounds,  are  precipi- 
tated. These  may  be  considered  as  formed  by  the  ammonolysis 
of  mercuric  salts  by  a  process  which  is  closely  analogous  to  the 
formation  of  an  ordinary  basic  (aquobasic)  salt  by  hydrolysis. 
Thus  the  partial  hydrolysis  of  mercuric  chloride  may  give  in 
solution : 

HgCl2  +  H.OH  ^±  H-O— HgCl  +  HC1 

Aquobasic 
Mercuric  Chloride 

In  the  presence  of  ammonia,  by  an  exactly  analogous  reaction, 
we  should  have : 

HgCl2  +  H.NH2  ^  H2N— HgCl  +  HC1 

Ammonobasic 
Mercuric  Chloride 

or          Hg(N03)2  +  H.NH2  ^±  H2N— HgNO3  +  HNO3 

Ammonobasic 
Mercuric  Nitrate 

The  hydrochloric  or  nitric  acid  would,  of  course,  unite  with 
the  excess  of  ammonia  present  to  form  ammonium  chloride, 
NH4C1,  or  ammonium  nitrate,  NH^.NOs.  Salts  of  many  other 
metals  undergo  ammonolysis  in  solutions  in  anhydrous  ammo- 
nia, but  the  ammonobasic  compounds  which  are  formed  are 
decomposed  by  water  in  almost  all  cases,  while  the  ammono- 
basic mercuric  compounds  are  stable  in  the  presence  of  water, 
either  because  of  their  extreme  insolubility  or  because  of  some 
specific  affinity  between  mercury  and  nitrogen.  (See  E.  C. 
Franklin,  J.  Am.  Chem.  Soc.  29,  35  (1907) ;  Am.  Chem.  J.  47, 
363  (1912) ). 


MAGNESIUM,  ZINC,   CADMIUM  AND  MERCURY     493 

Mercurous  salts  react  with  ammonia  as  though  they  were 
mixtures  of  a  mercuric  salt  with  mercury : 

Hg2Cl2  +  2  H.NH2  ^±  H2N.HgCl  +  Hg  +  NH4C1. 

The  metallic  mercury  colors  the  precipitate  formed  from 
mercurous  salts  black. 

If  ammonobasic  mercuric  chloride  is  dissolved  in  a  solution 
of  ammonium  chloride  in  anhydrous  ammonia,  the  ammonolysis 
may  be  reversed  exactly  as  the  hydrolysis  of  a  salt  may  be 
reversed  by  hydrochloric  acid  : 

H2N— Hg— Cl  +  NH4C1  ^±  HgCl2  +  2  NH3 

From  such  a  solution  a  compound  of  the  formula  HgCl2.2  NHa, 
which  contains  ammonia  of  crystallization  and  is  closely  analo- 
gous to  the  hydrates  of  other  salts  (p.  82),  may  be  crystallized. 

Nessler's  Reagent.  Mercuric  iodide,  HgI2,  which  is  almost 
wholly  insoluble  in  water,  dissolves  easily  in  a  solution  of  potas- 
sium iodide,  owing  to  the  formation  of  a  complex  salt,  potassium 
mercuric  iodide,  K2HgI4.  In  such  a  solution  sodium  hydroxide 
will  give  no  precipitate ;  but  if  ammonia  or  an  ammonium  salt  is 
added  to  the  alkaline  solution,  a  precipitate  of  ammonobasic- 
aquobasic-mercuric  iodide,  HO — Hg — NH — Hg — I,  is  formed. 
In  very  dilute  solutions  of  ammonia  the  solution,  which  is  called 
"  Nessler's  reagent,"  produces  a  brown  coloration  which  is  used 
for  the  detection  and  quantitative  estimation  of  ammonia. 

EXERCISES 

1.  How  much  crystallized  hydrate  of  magnesium  chloride  and  how 
much  ammonium  chloride  will  be  required  to  furnish  one  pound  (453 
grams)  of  anhydrous  magnesium  chloride  ? 

2.  How  many  liters  of  carbon  dioxide  at  20°  and  760  nun.  will  be 
given  by  heating  84  grains  of  magnesium  carbonate  ? 

3.  How  much  dolomite  would  be  required  to  give  a  pound  of  Epsom 
salts  ? 

4.  What  volume  of  gases  will  0.284  gram  of  mercuric  fulminate  give 
by  its  explosion,  supposing  the  temperature  of  the  gases  to  be  546°  ? 


CHAPTER  XXVIII 

METALS   OF   GROUP   III.     ALUMINIUM   FAMILY.     RARE 
EARTH   METALS 

WHILE  all  of  the  elements  of  both  divisions  of  Group  II  are  fully 
metallic  in  character  and  all  except  radium  are  comparatively 
common  and  their  compounds  well  known,  the  first  element, 
boron,  of  Group  III,  is  decidedly  nonmetallic,  and  aluminium  is  the 
only  metal  of  the  group  which  can  be  considered  very  common. 

Aluminium,  Al,  27.1,  is  found  in  a  great  variety  of  natural 
silicates,  especially  in  the  feldspar  and  mica  of  the  granites  and 
similar  rocks,  which  are  still  abundant  and  which  must  have 
been  much  more  common  in  early  geologic  time.  By  the  pro- 
longed action  of  water  and  the  forces  of  nature  such  rocks  have 
been  slowly  disintegrated.  The  potassium  and  sodium  of  the 
minerals,  have  been  partly,  though  by  no  means  completely, 
dissolved  and  removed,  and  a  hydrated  silicate  of  aluminium, 
mixed  with  fragments  of  quartz  and  of  partially  decomposed 
minerals,  has  been  left  in  an  extremely  fine  state  of  division. 
This  material,  after  transportation  for  some  distance  by  water, 
has  been  deposited  by  sedimentation  and  has  formed  immense 
beds  of  shales,  clays  and  soils.  Such  clays  and  shales  may  be 
considered  as  ores  of  aluminium,  though  they  contain  only  from 
15  to  30  per  cent  of  the  metal.  Pure  kaolin,  the  mineral  basis  of 
clay,  has  the  formula  Al2Si2O7.2H2O.  Aluminium  also  occurs  as 
the  oxide,  A12C>3,  in  crystals  known  as  ruby  and  sapphire,  which 
are  used  as  gems,  and  in  a  massive,  very  hard  form,  called  emery 
and  used  as  an  abrasive.  The  mineralogical  name  of  the  oxide 
is  corundum.  Bauxite,  the  hydrate,  A12O3.2  H2O,  usually 
containing  a  considerable  amount  of  the  hydrate  of  ferric  oxide, 
2  Fe2O3.3  H2O,  is  the  chief  source  from  which  aluminium  oxide 

494 


ALUMINIUM  495 

is  prepared  for  the  manufacture  of  the  metal.  Cryolite,  NaaAlF6, 
is  a  soft,  easily  fusible  mineral  found  in  large  quantities  in  Green- 
land, but,  so  far  as  known,  nowhere  else,  except  as  a  rare  mineral. 

Metallurgy.  Aluminium  was  first  prepared  by  the  German 
chemist  Wb'hler  in  1828  by  the  action  of  potassium  on  aluminium 
chloride,  but  he  obtained  only  a  very  small  quantity  in  the  form 
of  a  gray  powder.  Twenty-six  years  later  Sainte-Claire-Deville 
exhibited  in  Paris  a  quantity  of  the  metal,  which  he  obtained  by 
the  action  of  sodium  on  the  chloride,  and  the  element  aroused  a 
great  deal  of  interest  as  "  silver  from  clay."  The  desire  of  ob- 
taining the  metal  in  larger  quantities  led  to  the  development  of 
cheaper  and  better  methods  for  the  manufacture  of  sodium,  but 
as  the  valence  of  sodium  is  one  while  that  of  aluminium  is  three, 
and  the  atomic  weights  are  not  far  different,  it  must  always  take 
about  three  pounds  of  sodium  to  give  one  pound  of  aluminium, 
and  the  metal  manufactured  by  that  process  was  never  put  on 
the  market  at  a  price  below  $10  to  $12  a  pound. 

In  1885  Professor  Mabery  of  Cleveland,  at  a  meeting  of  the 
American  Association  for  the  Advancement  of  Science  held  in 
Ann  Arbor,  gave  an  account  of  a  new  electric  furnace  devised 
in  1882,  by  the  Cowles  brothers,  for  the  production  of  aluminium 
bronze.  They  had  discovered  that,  at  the  high  temperature 
of  the  electric  arc,  aluminium  oxide,  A^Os,  can  be  reduced  by 
carbon  to  the  metallic  form,  and  that  if  copper  is  present  the  alloy, 
aluminium  bronze,  can  be  obtained.  This  seems  to  have  been 
the  first  application  of  the  electric  furnace  to  an  industrial  pro- 
cess. Its  use  for  this  particular  purpose  was  short-lived.  A 
few  years  later  another  American,  C.  M.  Hall,  discovered  that 
aluminium  oxide  dissolves  easily  in  melted  cryolite,  and  that  if 
an  electric  current  is  passed  from  a  carbon  anode  through  the 
molten  mass  contained  in  an  iron  pot,  aluminium  is  deposited 
in  the  bottom  of  the  pot,  while  oxygen  liberated  at  the  anode 
combines  with  the  carbon  and  escapes  as  carbon  dioxide  (Fig. 
103).  Other  materials  are  now  used  wholly  or  in  part  in  place 
of  the  cryolite,  but  the  principles  used  in  the  process  are  not 
changed.  The  aluminium  oxide  is  obtained  by  heating  bauxite 


496 


A  TEXTBOOK  OF  CHEMISTRY 


with  carbon  in  an  electric  furnace.     The  iron,  silicon  and  other 

elements  in  the  bauxite  are  reduced  by  this  process  and  may 

be  separated  from  the 
fused,  pure  aluminium 
oxide. 

*  Thus  far  the  alu- 
minium oxide  used  for 
the  production  of  alu- 
minium has  not  been 
prepared,  commercially, 
from  clay,  but  has 
usually  been  made  from 
bauxite.  Very  recently 
Alfred  H.  Cowles  has 
Fig.  103  developed  a  process  by 

which  clay  mixed  with 

salt  and  charcoal  is  heated  in  a  current  of  air  and  steam,  giving 

the  reaction : 

Al2Si2O7  +  4  NaCl  +  2  H2O  =  2  Na2O.2  SiO2.  A12O3  +  4  HC1 

The  carbon  burns  to  carbon  monoxide  and  serves  to  render  the 
material  porous  and  easily  accessible  to  the  steam  and  air.  Iron, 
which  is  present,  volatilizes  as  ferric  chloride,  FeCl3. 

If  the  mixture  of  sodium  silicate  and  aluminate  is  mixed  with 
lime  and  heated,  an  insoluble  calcium  silicate  and  soluble  sodium 
aluminate  are  formed : 

2  Na2O.2  Si02.Al2O3  +  4  CaO 

=  2  Ca2SiO4  +  NaAlO2  +  Na3AlO3 

Calcium  &    •, .         . , 

Silicate  Sodmm  Alummate 

From  the  solution  of  sodium  aluminate,  aluminium  hydroxide 
may  be  precipitated  by  carbon  dioxide  : 

2  NaA102  +  C02  +  3  H2O  =  2  A1(OH)3  +  Na2CO3 

The  process  seems  promising  because  it  gives  three  valuable 
compounds,  hydrochloric  acid,  sodium  carbonate  and  aluminium 


ALUMINIUM:  THERMITE  497 

hydroxide,  with  the  use  of  cheap  raw  materials  (Journal  of 
Industrial  and  Engineering  Chemistry,  5,  331). 

Properties  of  Aluminium.  Aluminium  melts  at  658.7°  and 
boils  at  1800°.  It  has  a  specific  gravity  of  only  2.6,  almost  the 
same  as  that  of  glass  and  scarcely  more  than  one  third  that  of 
iron.  This  makes  it  useful  for  the  construction  of  apparatus 
which  should  be  light.  It  does  not  tarnish  readily  and  is  used 
to  some  extent  for  cooking  utensils.  It  is  not  attacked  by  water, 
even  at  the  boiling  point,  but  dissolves  readily  in  alkalies  or  in 
acids,  forming  aluminates  with  the  alkalies,  such  as  NaAlO2,  and 
salts  with  acids,  such  as  A^SOJs.  Aluminium  is  rather  easily 
corroded  by  salt  solutions.  Aluminium  which  has  been  amalga- 
mated by  bringing  it  into  contact  with  a  dilute  solution  of  mer- 
curic chloride  becomes  active  and  will  decompose  water  rapidly 
at  ordinary  temperatures.  In  this  condition  it  is  in  very  sharp 
contrast  with  amalgamated  zinc,  which  is  not  attacked  by  hy- 
drochloric or  sulfuric  acid  because  of  its  homogeneous  surface. 
The  surface  of  amalgamated  aluminium  is  gray  and  evidently 
nonhomogeneous.  Aluminium  is  used  to  some  extent  for  elec- 
tric conductors  in  place  of  copper.  It  is  often  used  as  an  addi- 
tion to  cast  iron,  greatly  improving  its  quality. 

Alloys.  The  best  known  alloy  is  aluminium  bronze,  composed 
of  copper  with  5-12  per  cent  of  aluminium.  It  resembles  gold 
very  closely  in  appearance  and  does  not  tarnish  readily.  Magna- 
lium,  an  alloy  with  a  small  amount  of  magnesium,  is  very  light 
and  is  much  more  easily  worked  on  a  lathe  than  aluminium 
itself.  Alloys  containing  from  2  to  10  per  cent  of  copper  are 
used  for  castings  for  automobiles  and  for  other  purposes  where 
lightness  is  desirable. 

Goldschmidt's  Thermite  Process.  Aluminium  has  a  very 
strong  affinity  for  oxygen,  as  shown  by  the  difficulty  with  which 
it  is  reduced.  The  heat  of  combustion  of  aluminium  is  : 

2  Al  +  3  O  =  A12O3  +  380,000  calories 
That  of  iron,  if  it  could  be  burned  to  ferric  oxide,  is  : 
2  Fe  +  3  O  =  Fe2O3  +  195,000  calories 


498  A  TEXTBOOK  OF   CHEMISTRY 

From  these  values  it  is  evident  that  the  reaction 
Fe203  +  2  Al  =  A12O3  +  2  Fe 

may  occur  with  the  evolution  of  a  large  amount  of  heat.  Gold- 
schmidt  has  made  use  of  this  principle  for  the  production  of  very 
high  temperatures  and  also  for  the  reduction  of  chromic  oxide 
and  other  refractory  oxides.  Because  the  aluminium  oxide  is 
not  volatile,  the  heat  of  the  reaction  is  not  dissipated  by  the  for- 
mation of  a  vapor,  and  a  temperature  high  enough  to  melt  iron 
or  steel  may  be  easily  obtained  for  the  welding  of  steel  rails,  per- 
foration of  iron  plates  and  similar  purposes. 

The  thermite  process  is  also  very  useful  for  the  production  of 
chromium  and  other  metals  which  it  is  difficult  to  obtain  in 
other  ways.  A  special  advantage  of  the  process,  in  some  cases, 
is  that  the  metals  obtained  in  this  way  are  free  from  carbon. 

Aluminium  Chloride,  A1C13,  is  easily  prepared  by  heating 
aluminium  turnings  in  chlorine  or  in  hydrochloric  acid.  It 
sublimes  at  183°  under  atmospheric  pressure  and  melts  at  193° 
under  increased  pressure.  Aluminium  chloride  dissolves  in 
water  with  the  evolution  of  considerable  heat.  If  the  solution 
is  not  too  dilute,  the  addition  of  concentrated  hydrochloric  acid 
will  cause  a  crystalline  hydrate,  A1C13.H2O,  to  separate.  On 
heating  this  hydrate  it  loses  hydrochloric  acid  and  aluminium 
oxide  is  left : 

2  Aids  +  3  H2O  =  A12O3  +  6  HC1 

Aluminium  chloride  forms  with  many  organic  substances 
extremely  reactive  addition  compounds  which  are  used  in  a 
variety  of  syntheses.  For  these  reactions  water  must  be  care- 
fully excluded,  and  the  hydrate  of  the  chloride  cannot  be  used 
at  all  for  such  purposes.  All  of  these  facts  indicate  very  clearly 
that  the  water  of  the  hydrate  is  in  a  state  of  intimate  chemical 
combination  such  as  to  greatly  modify  the  relation  which  the 
chlorine  and  aluminium  have  in  the  anhydrous  chloride.  Solu- 
tions of  aluminium  chloride  are  partially  hydrolyzed  by  water, 
and  react  strongly  acid.  Nearly  all  salts  of  trivalent  or  quadri- 
valent metals  act  in  the  same  way. 


ALUMINIUM  COMPOUNDS  499 

Aluminium  Fluoride,  A1F3.  Either  aluminium  or  aluminium 
hydroxide  dissolves  easily  in  an  aqueous  solution  of  hydrofluoric 
acid,  forming  a  supersaturated  solution  from  which  anhydrous 
aluminium  fluoride  slowly  separates  in  small  crystals.  The 
double  fluoride,  Na3AlFe,  is  found  in  nature  as  the  mineral 
cryolite,  and  was  formerly  used  to  dissolve  the  oxide  for  the 
electrolytic  preparation  of  metallic  aluminium. 

Aluminium  Hydroxide,  A1(OH)3,  is  precipitated  from  solutions 
of  aluminium  salts  on  the  addition  of  an  alkaline  hydroxide. 
It  forms  a  voluminous,  gelatinous  precipitate  which  dissolves 
either  in  solutions  of  strong  acids  or  strong  bases.  For  this 
reason  it  is  called  amphoteric,  meaning  that  it  has  both  basic  and 
acid  properties.  It  seems  probable  that  in  contact  with  an  acid 
its  hydroxyl  combines  with  the  hydrogen  of  the  acid,  while  in 
contact  with  a  base  its  hydrogen  combines  with  the  hydroxyl  of 
the  base.  The  compound  may  be  considered,  therefore,  either 
as  a  triacid  base  or  as  a  tribasic  acid.  It  is,  of  course,  very  weak 
both  as  an  acid  and  as  a  base.  Salts  in  which  it  is  the  cation  are 
hydrolyzed  in  solution  and  have  an  acid  reaction  : 

A12(SO4)3  +  6  HOH  ^T  2  A1(OH)8  +  3  H2SO4 

Those  salts  in  which  the  aluminium  forms  part  of  the  anion  are 
also  hydrolyzed  and  have  an  alkaline  reaction  : 

Na3AlO3  +  3  HOH  ^±  A1(OH)8  +  3  NaOH 

When  heated,  aluminium  hydroxide  loses  water  and  is  con- 
verted into  the  oxide,  A12O3.  It  conducts  itself  very  much  as 
the  silicic  acids,  however,  losing  a  part  of  the  water  very  easily 
at  ordinary  temperatures  but  requiring  ignition  at  bright  redness 
to  expel  the  last  portions. 

Aluminium  Oxide,  A12O3,  is  found  in  nature  as  the  mineral 
corundum.  In  its  massive,  not  very  pure  forms,  it  is  called 
emery.  Next  to  the  diamond  it  is  the  hardest  mineral  known 
and  has  been  long  used  as  an  abrasive  for  grinding  and  polish- 
ing glass  and  metals.  It  has  now  been  partly  displaced  for  these 
uses  by  carborundum,  SiC,  which  is  much  harder.  Crystalline 


500  A  TEXTBOOK  OF  CHEMISTRY 

forms  of  corundum  colored  blue  by  some  foreign  substance  are 
called  sapphires,  or  other  forms  colored  red  by  chromium  are 
called  rubies.  The  latter  are  now  made  artificially.  A  fused 
oxide,  prepared  in  the  electric  furnace,  is  used,  under  the  name 
of  alundum,  as  an  abrasive  and  also  as  a  refractory  material. 

Ignited  aluminium  oxide  is  insoluble  in  acids  but  may  be 
brought  into  solution  slowly  by  fusion  with  sodium  pyrosulfate, 
Na2S2O7.  Aluminium  oxide  is  reduced  by  carbon  at  the  tem- 
perature of  the  electric  furnace,  but  cannot  be  reduced  at  lower 
temperatures.  Before  metallic  aluminium  was  prepared  by 
electrical  processes,  the  anhydrous  aluminium  chloride  used  for 
the  preparation  of  the  metal  was  obtained  by  heating  a  mixture 
of  the  oxide  with  carbon  in  a  current  of  chlorine : 

A12O3  +  3  C  +  3  C12  =  2  A1C18  +  3  CO 

Aluminium  Sulfate,  A12(SO4)3.18  H2O,  is  prepared  by  the 
decomposition  of  clay  with  sulfuric  acid.  A  more  or  less  pure 
sulfate  containing  some  ferric  sulfate  is  extensively  used  under 
the  name  of  "  alum  "  for  the  clarification  and  purification  of 
water.  If  a  solution  of  the  sulfate  is  mixed  with  a  water  con- 
taining calcium  bicarbonate,  an  insoluble,  gelatinous  precipitate 
of  aluminium  hydroxide  is  formed  : 

A12(SO4)3  +  3  CaH2(C03)2  =  2  A1(OH)8  +  3  CaSO4  +  3  CO2 

The  precipitate  collects  fine  particles  of  clay  and  also  bacteria 
which  are  suspended  in  the  water,  and  by  suitable  filtration  clear 
water,  nearly  or  quite  free  from  disease  germs,  is  obtained.  The 
amount  of  aluminium  sulfate  required  is  so  small  that  the  calcium 
sulfate  formed  does  not  seriously  increase  the  permanent  hard- 
ness of  the  water.  All  of  the  aluminium  added  is  removed  in 
the  filtration. 

Alums,  M'M'"(SO4)2.12H2O.  By  adding  potassium  sul- 
fate, K2SO4,  to  a  solution  of  aluminium  sulfate  a  compound 
having  the  composition  KA1(SO4)2.12H2O,  and  known  since 
early  times  under  the  name  of  alum,  is  formed.  It  crystallizes 
in  octahedra,  which,  with  care,  may  be  obtained  in  very  perfect 
forms  of  large  size. 


ALUMS,  EARTHENWARE  501 

Alum  was  formerly  much  used  as  a  mordant  in  dyeing.  The 
aluminium  hydroxide  formed  by  its  hydrolysis  combines  with 
many  coloring  matters  to  form  insoluble  compounds  called 
lakes.  These  compounds  attach  themselves  strongly  to  the 
fibers  of  the  cloth  and  cannot  be  removed  by  washing.  Alu- 
minium sulf ate  and  the  aluminates  have  largely  displaced  alum 
for  such  uses  because  the  potassium  sulfate  is  expensive  and 
unnecessary. 

The  potassium  of  alum  may  be  replaced  by  ammonium  or 
other  univalent  metals,  the  aluminium  may  be  replaced  by  ferric 
iron  or  other  trivalent  metals  and  even  the  sulfate  radical,  SO4, 
may  be  replaced  by  the  selenate  radical,  SeO4.  This  gives 
a  great  variety  of  alums,  all  of  which  crystallize  in  octahedra 
and  are  isomorphous.  A  crystal  of  any  alum  will  grow  in  a 
supersaturated  solution  of  any  other. 

The  following  may  be  given  as  illustrations  of  the  alums : 

Ammonium  Alum  NH4A1(SO4)2 .  12  H2O 

Ammonium  Ferric  Alum  NH4Fe(SO4)2 .  12  H2O 

Chrome  Alum  KCr(SO4)2 .  12  H2O 

Rubidium  Alum  RbAl(SO4)2 .  12  H2O 

Brick,  Earthenware,  Porcelain.  Aluminium  silicate  melts 
only  at  very  high  temperatures,  but  the  presence  of  other  com- 
mon metals,  such  as  iron,  calcium,  magnesium,  sodium  or  potas- 
sium, lowers  the  melting  point.  Ordinary  clays  contain  com- 
pounds of  these  metals  distributed  as  very  fine  particles  through- 
out their  mass,  and  when  such  clays  are  heated  to  a  high 
temperature,  the  particles  melt  and  cause  the  material  to  sinter 
together  to  a  strong  but  very  porous  mass.  In  addition  to  this 
property  of  sintering  without  fully  melting,  the  original  clays 
become  plastic  when  mixed  with  water,  and  in  this  condition 
may  be  molded  into  bricks  or  into  the  "  biscuit  "  forms  which 
furnish  the  basis  of  earthenware  or  porcelain.  The  plasticity 
of  the  clay  seems  to  be  closely  connected  with  its  colloidal 
character. 

In  order  to  give  them  a  surface  which  is  smoother  and  im- 


502  A  TEXTBOOK  OF  CHEMISTRY 

pervious  to  water,  articles  of  earthenware  and  porcelain  must 
be  covered  with  a  glaze.  Several  methods  of  glazing  are  in 
use.  One  method  is  to  throw  salt  into  the  furnace  after  the 
biscuit  has  been  well  burned.  The  sodium  chloride  reacts  with 
the  silica  and  alumina  of  the  clay  and  the  moisture  of  the  air 
to  form  hydrochloric  acid,  which  escapes,  and  a  fusible  silicate 
is  formed,  which  melts  and  covers  the  surface  with  a  glass. 
Other  glazes  are  made  from  mixtures  containing  lead  oxide. 
Some  glazes  of  this  type  are  not  wholly  insoluble  in  water,  and 
England  has  enacted  stringent  laws  requiring  lead  glazes  to  be 
highly  insoluble  when  the  articles  are  to  be  used  to  contain  or 
cook  food.  Porcelains  are  usually  glazed  by  the  application  of 
finely  powdered  feldspar  and  subjecting  them  to  a  temperature 
which  causes  it  to  melt  and  run  into  the  surface. 

Ultramarine.  Small  quantities  of  a  beautiful  blue  stone 
called  lapis  lazuli  are  found  in  nature.  When  powdered  this 
stone  gives  a  beautiful  blue  pigment  which  is  not  affected  even 
by  long  exposure  to  the  light.  The  mineral  is  so  rare,  however, 
that  during  the  first  years  of  the  nineteenth  century  the  pig- 
ment was  sold  to  artists  at  $60  an  ounce.  In  1828  Gmelin  dis- 
covered that  the  material  can  be  made  artificially  by  heating 
mixtures  of  clay,  sodium  sulfate,  charcoal  and  sulfur.  The 
artificial  product  is  fully  equal  to  the  natural  mineral  for  the 
uses  to  which  it  is  applied  and  is  now  sold  at  a  few  cents  per 
pound.  By  changing  the  method  of  manufacture,  other  com- 
pounds, ranging  in  color  from  reddish  violet  to  bluish  green, 
are  also  made.  In  spite  of  a  very  large  amount  of  work  devoted 
to  the  preparation  and  analysis  of  these  compounds  it  has  not 
been  possible  to  assign  definite  formulas  to  them. 

The  Rare  Earths.  These  include  the  oxides  of  a  number  of 
elements  all  of  which  are  characterized  by  being  trivalent  and 
by  forming  oxalates  which  are  insoluble  in  dilute  mineral  acids. 
They  resemble  each  other  very  closely  in  all  their  properties 
and  in  the  types  of  compounds  which  they  form.  Their  salts 
are  isomorphous  and  do  not  as  a  rule  differ  greatly  in  solubility. 
Because  of  these  slight  differences  they  cannot,  with  the  possible 


ALUMINIUM  FAMILY.    RARE   EARTH  METALS     503 

exception  of  cerium,  be  separated  from  each  other  quantita- 
tively, and  their  preparation  in  a  pure  state  requires  a  long-con- 
tinued series  of  fractional  crystallizations  or  precipitations.  In 
fact,  some  of  them  have  not  been  obtained  in  a  high  state  of 
purity.  These  elements  may  be  divided  into  two  groups  :  the 
Cerium  group,  including  cerium,  lanthanum,  neodymium,  pra- 
seodymium, samarium,  europium  and  gadolinium ;  and  the 
Yttrium  group,  including  terbium,  dysprosium,  holmium, 
yttrium,  erbium,  thulium,  ytterbium,  scandium  and  lutecium. 
The  members  of  the  cerium  group  may  be  separated  roughly 
from  those  of  the  yttrium  group  by  making  use  of  the  fact  that 
the  elements  of  the  former  form  double  sulfates  with  sodium, 
sulfates  which  are  insoluble  in  a  saturated  solution  of  sodium 
sulfate,  while  the  corresponding  compounds  of  the  yttrium 
group  are  soluble. 

*  Scandium,  Sc,  44.1.    When  Mendeleeff  proposed  the  Periodic 
System  of  the  elements  in  1869,  he  predicted  that  several  ele- 
ments not  then  known  would  probably  be  discovered  in  the 
future.     Among  these  was  an  element  which  he  called  "  eka- 
boron "  which  should  have  an  atomic  weight  of  about  44  and 
form  compounds  similar  to  those  of  aluminium.     Ten  years 
later   Nilson   found   scandium   among  the  elements  found  in 
gadolinite  and  euxenite,  and  shortly  after  Mendeleeff  pointed 
out  that  this  is  in  reality  the  "  ekaboron  "  which  he  had  pre- 
dicted.    The     hydroxide,     Sc(OH)3,     oxide,     Sc2O3,     sulfate, 
Sc2(SO4)3.6H2O,   oxalate,    Sc2(C2O4)3.6  H2O    and  other    salts 
are  known. 

*  Yttrium,  Y,  89,  is  found  in  gadolinite,  xenotime  and  monazite. 
Its  compounds  resemble  those  of  scandium.     The  oxide,  Y2O3, 
chloride,  YC13,  sulfide,  Y2S3,  phosphate,  YPO4,  and  bromate, 
Y(BrO3)3.9H2O,  may  be  mentioned. 

*  Lanthanum,  La,  139,  is  the  most  positive  of  the  rare  earth 
metals.      Its  oxide,  La^Os,  combines  with  water,  much  as  lime 
does,  forming  the  hydroxide,  La(OH)3,  which  turns  litmus  paper 
blue.     The    hydroxide    also    absorbs    carbon    dioxide    from 
the    air,    forming    the    carbonate,  La2(CO3)3.     The    oxalate, 


504  A  TEXTBOOK  OF  CHEMISTRY 

La2(C2O4)3.9  H2O,  is  difficultly  soluble,  as  are  the  oxalates  of 
all  of  the  rare  earth  metals. 

*  Ytterbium,   Yb,   172,  gives  the  sulfate,  Yb2(SO4)3.8  H2O, 
carbonate,     Yb2(CO3)3.4  H2O,    acetate,   Yb(C2H3O2)3.4  H2O, 
oxalate,  Yb2(C2O4)3.10  H2O,  and  many  other  salts. 

*  Praseodymium,   Pr,    140.6,   and   Neodymium,   Nd,    144.3. 
In  1842  Mosander  obtained  from  cerite  (a  silicate  containing 
cerium,  lanthanum,  praseodymium  and  neodymium)  an  oxide 
of  a  metal  to  which  he  gave  the  name  didymium.     In  1885 
Auer.  v.  Welsbach  discovered  that  by  a  long  series  of  crystalli- 
zations the  double  nitrate  of  the  metal  previously  called  didym- 
ium could  be  separated   into  two   compounds,  praseodymium 
ammonium  nitrate,  Pr2(NO3)3.2  NH4NO3.4  H2O,  and  neodym- 
ium ammonium  nitrate,  Nd(NO3)3.2 NH4NO3.4H2O.     These 
two  salts  are  isomorphous  and  do  not  differ  very  greatly  in 
solubility.     It  was  necessary,  therefore,  to  repeat  the  crystalli- 
zation many  hundreds  of  times  in  such  a  manner  that  the  more 
soluble  portions  were  transferred  in  one  direction  through  the 
crystallizing  dishes  and  the  less  soluble  portions  in  the  other 
direction,  while  portions  of  the  same  degree  of  separation  were 
systematically  united.     This  process  of  fractional  crystalliza- 
tion has  been  much  used  in  the  separation  of  the  metals  of  the 
rare  earths.     (See  Auer.  v.  Welsbach,  Monatshefte  fur  Chemie, 
6,  477 ;  Baxter  and  Chapin,  J.  Amer.  Chem.  Soc.  33,  5 ;  James, 
ibid.  30,  182;  31,  913.) 

Typical  compounds  of  praseodymium  and  neodymium  are, 
Pr2O3,  PrO2,  Nd2O3,  Pr2(SO4)3.8  H2O,  Nd2(SO4)3.8  H2O, 
Pr2(C2O4)3 . 10  H2O,  Nd2(C2O4)3 . 10  H2O,  Pr(BrO3)3 . 9  H2O, 
Nd(Br03)3.9H20. 

The  oxalates  of  nearly  all  of  the  rare  earth  metals  are  very 
difficultly  soluble  and  are  often  used  as  a  means  of  separating 
these  metals  from  those  of  other  groups. 

The  salts  of  praseodymium  are  green  in  color,  those  of 
neodymium  are  rose-colored,  the  two  colors  being  comple- 
mentary very  much  as  those  of  the  salts  of  cobalt  and  nickel 
are. 


ALUMINIUM  FAMILY.    RARE  EARTH  METALS     505 

*  Samarium,  Sm,  150.4,  is  found  in  samarskite,  a  columbate  of 
metals  of  the  rare  earths,  and  was  first  partially  separated  from 
the  "  didymium  "  of  that  mineral  by  Lecoq  de  Boisbaudran. 
It  is  less  basic  than  praseodymium  and  neodymium,  and  its 
double  nitrate  with  magnesium  is  more  easily  decomposed  by 
heat,   a  method  sometimes  used  in  separations.     The  oxide, 
Sm2O3,  and  solutions  of  its  salts  are  yellow.     It  forms  a  chloride, 
SmCl2,  in  which  the  metal  is  bivalent,  but  in  nearly  all  of  its 
salts    it    is    trivalent.     A  considerable  number  of  salts  have 
been  prepared,    such   as   samarium  sulfate,  Sn^SO^s-S  H2O, 
the        nitrate,       Sm(NO3)3.6H2O       and       the       carbonate, 
Sm2(CO3)3.3H2O. 

*  Europium,  Eu,  152,  Gadolinium,  Gd,  157.3,  and  Terbium, 
Tb,  159.2,  form  a  group   of  weakly  basic  earths  intermediate 
between  the  cerium  earths  on  the  one  side  and  the  yttrium 
earths  on  the  other.     They  have  been  separated  by  tedious 
fractional  crystallizations  and  finally  identified  by  means  of 
their    spectra    and    determinations    of    their    atomic    weights. 
Europium  oxide,  Eu2O3,  has  a  light  rose  color,  and  the  sulfate, 
Eu2(SO4)3.8H2O,     is     also    rose-colored.      Gadolinium    oxide, 
Gd2O3,   and    sulfate,    Gd2(SO4)3.8H2O,  are  white.      Terbium 
oxide,  Tb2Os,  is  also  white ;  but  a  higher  oxide,  possibly  TbO2, 
but  not  yet  obtained  pure,  is  dark  brown  or  black  according  to 
the  method  of  preparation. 

*  Holmium,  Ho,  163.5,  is  obtained  from  euxenite,  a  columbate 
and  titanate  of  the  yttrium  and  cerium  earths.     The  solubility 
of  the  double  sulfate  with  ammonium  seems  to  lie  between 
those  of  yttrium  and  erbium  (Holmberg).      Apparently  pure 
compounds  have  not  yet  been  prepared.     The  metal  is  named 
from  Stockholm. 

*  Dysprosium,   Dy,   162.5,  has  been  obtained  in  its  purest 
form  by  fractional  crystallization  of  salts  of  ethyl  sulfuric  acid, 
H(C2H5)S04.     The    dysprosium    salt   is    Dy(C2H5SO4)3.     The 
oxide,  Dy2O3,  is  white.     The  bromate,   Dy(BrO3)3.9  H2O,  is 
yellow. 

*  Erbium,  Er,  167.7,  is  found  among  the  yttrium  earths  from 


506  A  TEXTBOOK  OF  CHEMISTRY 

euxenite  and  other  sources.     The  oxide,  Er203,  is  rose-colored, 
as  are  also  the  sulfate,  Er2(SO4)3.8  H2O,  and  other  salts. 

*  Thulium,  Tu,   168.5,  is  found  in  euxenite  and  other  rare 
earth  minerals.     It  seems  to  have  been  separated  in  the  purest 
condition   by  the  fractional   crystallization   of    the  bromates 
(James,  J.  Amer.  Chem.  Soc.,  33,  1332).     The  oxide,  Tu2O3,  is 
white,  with  a  faint  green  tint.     The  bromate,  Tu(BrO3)3. 9  H2O, 
separates  in  pale,  bluish  green  prisms,  isomorphous  with  the 
bromates  of  other  rare  earth  metals. 

*  Lutecium,  Lu,  174,  is  one  of  the  more  recently  discovered 
elements  of  this  group  and  has  been  separated  from  the  gado- 
linite  earths. 

*  Gallium,  Ga,  69.9,  was  also  predicted  by  Mendeleeff  under 
the  name  of  "  eka-aluminium."     Unlike  the  other  metals  of  the 
group  thus  far  described,  it  forms  two  classes   of  compounds, 
those  in  which  it  is  bivalent  and  others  in  which  it  is  trivalent. 
The  chlorides  are  GaCl2  and  GaCl3 ;  the  sulfates,  GaSO4  and 
Ga2(SO4)3.18  H2O.     The  former  is  oxidixed  by  potassium  perman- 
ganate as  ferrous  sulfate  is.     The  alum,  NH4Ga(SO4)2.12H2O, 
is  isomorphous  with  ordinary  alum. 

Indium,  In,  114.8,  was  discovered  by  means  of  the  blue  line 
of  its  spectrum  by  Reid  and  Richter  shortly  after  the  methods 
of  spectrum  analysis  had  been  developed  by  Bunsen  and  Kirchoff. 
The  analysis  of  its  oxide  gave  about  76  parts  of  indium  for  16 
parts  of  oxygen,  and  an  atomic  weight  of  76  was  at  first  assigned 
to  the  element.  But  Mendeleeff  pointed  out  that  this  would 
place  it  between  arsenic  and  selenium,  where  there  is  no  vacant 
place  in  the  Periodic  System,  and  also  that  its  properties  did  not 
agree  with  such  a  position  in  the  table.  He  suggested,  there- 
fore, that  the  formula  of  the  oxide  is  In2O3  and  the  atomic 
weight  114.  The  determination  of  the  specific  heat  gave  the 
value  0.056.  This  points  to  an  atomic  weight  of  6.2/0.056  =  110, 
which  agrees  fairly  well  with  an  atomic  weight  of  114  but  would 
not  agree  at  all  with  the  value  of  76.  The  preparation  of  an 
ammonium  alum,  NH4In(SO4)2.12H2O,  soon  after  this,  gave 
further  support  for  the  accepted  formula. 


ALUMINIUM  FAMILY.    RARE  EARTH  METALS     507 

Indium  is  a  soft,  white  metal,  which  melts  at  155°.  It  gives 
three  chlorides,  InCl,  InCl2,  InCl3,  but  the  compounds  in  which 
it  is  trivalent  are  most  stable  and  best  known. 

Thallium,  Tl,  204,  was  also  discovered  by  means  of  the  spec- 
troscope. Crookes  found  it  in  1861  in  the  slimes  from  sulfuric 
acid  made  at  Tilkerode  in  the  Harz.  He  named  it  thallium 
from  the  Latin  word  thallus,  meaning  a  young  twig,  because  of 
a  brilliant  green  line  in  its  spectrum.  Metallic  thallium  is  a 
bluish  white,  soft  metal,  somewhat  resembling  lead. 

Thallium  forms  thallous  compounds  in  which  it  is  univalent, 
and  thallic  compounds  in  which  it  is  trivalent.  Of  the  former, 
thallous  oxide,  T^O,  thallous  chloride,  T1C1,  thallous  hydroxide, 
T1OH.H2O,  and  thallous  sulfide,  T12S,  may  be  mentioned. 
The  last  is  a  black  precipitate  nearly  insoluble  in  acetic  acid 
but  soluble  in  mineral  acids.  Thallous  iodide,  Til,  is  also  very 
difficultly  soluble. 

Among  the  thallic  compounds  are  thallic  chloride,  T1C13.H2O, 
thallic  nitrate,  T1(NO3)3.8  H2O,  and  thallic  sulfide,  T12S3. 

EXERCISE 

Assuming  the  specific  heat  of  aluminium  oxide  as  0.217,  that  of  iron  as 
0.15,  and  the  heat  of  fusion  of  iron  as  23  calories  per  kilogram,  what  is 
the  maximum  temperature  which  could  be  reached  by  the  reaction  of  a 
thermite  consisting  of  ferric  oxide  and  metallic  aluminium  ? 


CHAPTER  XXIX 
TIN  AND   LEAD 

IT  has  been  pointed  out  (p.  361)  that  tin  and  lead  belong  to 
the  carbon  group  of  the  Periodic  System  and  that  each  gives 
an  oxide  (SnC>2  and  PbCy  resembling  carbon  dioxide,  CO2,  and 
silicon  dioxide,  SiO2,  in  formula  and  in  some  other  properties, 
especially  in  their  acidic  character.  Both  of  these  elements 
are  clearly  metals  rather  than  non-metals  in  most  of  their 
properties. 

Tin  (Sn,  119).  Occurrence,  Metallurgy.  Tin  is  rather  re- 
markable in  that,  although  it  is  an  element  which  is  found  in 
sufficient  quantity  so  that  it  is  a  common  metal  for  household 
and  commercial  use,  there  are  only  a  very  few  localities  in  the 
world  where  its  ores  can  be  profitably  mined.  One  of  the  oldest 
of  these  is  Cornwall  in  England,  where  tin  has  been  obtained 
for  nearly  or  quite  twenty  centuries  and  which  furnished  a  large 
part  of  the  tin  used  in  the  world  until  comparatively  modern 
times.  The  world's  supply  of  tin  for  one  year  is  approximately 
80,000  tons,  and  this  comes  almost  entirely  from  Banca  and  the 
East  India  islands,  Tasmania,  Bolivia  and  Cornwall.  The  only 
important  ore  is  cassiterite,  stannic  oxide,  SnO2.  This  is  found 
sometimes  in  veins,  sometimes  as  a  heavy  gravel,  called  "  stream 
tin."  The  metallurgy  is  comparatively  simple,  consisting  in  the 
reduction  of  the  oxide  by  means  of  charcoal  or  coal.  On  ac- 
count of  its  value,  the  recovery  of  tin  from  tin  scrap  has  also 
assumed  considerable  commercial  importance.  Several  methods 
are  in  use,  one  of  the  best  being  the  treatment  of  the  scrap  with 
dry  chlorine  gas,  which  converts  the  tin  into  stannic  chloride, 
SnCl4,  but  leaves  the  iron  comparatively  unattacked.  Stannic 
chloride  is  volatile  and  can  be  easily  separated. 

508 


TIN  509 

Uses  of  Tin.  Alloys.  Tin  Plate.  Tin  is  not  affected  by  dry 
or  moist  air  or  by  water,  even  at  the  boiling  point  or  higher. 
It  is,  for  this  reason,  the  most  suitable  of  the  cheaper  metals 
for  the  tubes  of  condensers  to  be  used  in  the  ^reparation  of 
distilled  water.  The  principal  use  of  the  metal  is  in  the  manu- 
facture of  tin  plate  —  sheet  iron  which  has  been  covered  with  a 
thin  coating  by  dipping  the  carefully  cleaned  metal  in  a  vat  of 
melted  tin.  Tin  is  more  electropositive  than  iron,  hence  when 
the  two  metals  are  in  contact  with  water  or  a  dilute  acid  the 
tendency  is  for  the  iron  to  corrode  while  the  tin  is  protected. 
For  this  reason  tin  vessels  rust  through  rapidly  as  soon  as  the 
iron  is  exposed  at  any  point  —  exactly  contrary  to  the  conduct 
of  iron  coated  with  zinc  (p.  482).  "  Terne  plate,"  which  is 
used  for  roofing  purposes,  is  covered  with  an  alloy  of  lead  and 
tin,  the  lead  being  used  because  it  is  very  much  cheaper  than  tin. 

Solder  is  an  alloy  of  lead  and  tin,  used  to  join  pieces  of  tin 
plate  in  making  culinary  vessels  of  all  kinds.  Common  solder 

(contains  equal  parts  of  the  metals,  but  fine  solder,  containing 
more  tin,  and  coarse  solder,  containing  more  lead,  are  often  used. 

Tin  was  formerly  much  used  in  gun  metal  and  bell  metal  and 
is  still  used  in  statuary  and  ornamental  bronzes,  which  are  alloys 
of  tin  and  copper,  usually  with  a  little  lead  and  zinc. 

Tin,  antimony  and  lead  are  the  principal  constituents  of 
Britannia  metal,  pewter  and  Babbitt  metal.  Copper  and  other 
metals  are  sometimes  added. 

Tin  melts  at  231.9°.  It  is  convenient  to  remember  the  melting 
points  of  tin  (232°),  lead  (327°)  and  zinc  (419°)  as  about  100° 
apart,  with  tin  the  lowest  and  zinc  the  highest.  Tin  boils  at  a 
high  temperature,  but  the  boiling  point  has  not  been  determined. 
The  specific  gravity  is  7.30. 

The  metal  oxidizes  slowly,  when  heated  above  its  melting 
point  in  the  air,  to  stannic  oxide,  SnO2.  It  dissolves  as  stannous 
chloride,  SnCl2,  in  concentrated  hydrochloric  acid  and  is  con- 
verted by  nitric  acid  into  a  mixture  of  stannic  and  metastannic 
acids,  SnO2.H2O,  which  is  insoluble  in  an  excess  of  the  acid  or 
in  water. 


510  A  TEXTBOOK  OF  CHEMISTRY 

Compounds  of  Tin.  Tin  forms  s  tan  nous  compounds,  such  as 
SnCl2,  in  which  it  is  bivalent,  and  stannic  compounds,  as  SnCU, 
in  which  it  is  quadrivalent.  In  the  former  it  is  rather  strongly 
basic  and  metallic,  in  the  latter  much  weaker  as  a  base  and  in 
some  of  the  compounds  distinctly  acidic. 

*  Stannous  Oxide,  SnO.  If  a  solution  of  potassium  carbonate, 
K2CO3,  is  added  to  a  solution  of  stannous  chloride,  SnCl2,  a 
white  precipitate  having  the  composition  2  SnO.H2O  is  formed ; 
but  this  loses  water  on  heating  the  solution,  especially  if  a 
little  alkali  is  present,  much  as  cupric  hydroxide,  Cu(OH)2,  does 
(p.  431),  and  is  changed  to  black  insoluble  stannous  oxide,  SnO. 
The  original  precipitate  is  amphoteric,  like  aluminium  hydroxide, 
Al(OH)s,  and  dissolves  either  in  acids  or  in  alkalies. 

Stannous  Chloride,  SnCl2.2  H2O,  is  easily  obtained  by  dis- 
solving tin  in  concentrated  hydrochloric  acid.  The  anhydrous 
salt  can  be  prepared  by  heating  this  hydrate  in  a  stream  of 
hydrochloric  acid.  It  boils  at  606°. 

Stannous  chloride  is  frequently  used  in  the  laboratory  as  a 
reducing  agent  because  of  its  strong  tendency  to  take  up  chlorine 
or  other  elements  and  pass  over  to  the  stannic  form.  A  solution  ' 
of  stannous  chloride  to  which  an  excess  of  sodium  hydroxide 
has  been  added,  forming  sodium  stannite,  NaHSnO2  or  Na2SnO2, 
is  also  a  powerful  reducing  agent. 

Stannous  chloride  and  several  other  compounds  of  tin  are  ex- 
tensively used  as  mordants  in  dyeing. 

Stannous  Sulphide,  SnS,  is  precipitated  from  acid  solutions  of 
stannous  salts  as  a  dark  brown,  almost  black  compound.  It 
does  not  dissolve  in  colorless  ammonium  sulfide,  (NH4)2S,  but 
the  yellow  ammonium  or  sodium  sulfides,  which  contain  poly- 
sulfides,  (NH4)2S2,  etc.,  dissolve  it  as  ammonium  or  sodium 
sulfostannate,  (NH4)2SnS3  or  Na2SnS3.  From  such  a  solution 
acids  precipitate  yellow  stannic  sulfide,  SnS2. 

Stannic  Oxide,  SnO2,  is  prepared  by  the  oxidation  of  tin  in 
air  at  a  high  temperature  or  by  treating  tin  with  nitric  acid  and 
igniting  the  mixture  of  stannic  acids  which  is  formed.  It  is 
also  found  as  the  crystalline  mineral,  cassiterite,  in  nature. 


STANNIC  ACIDS 


511 


Stannic  oxide  does  not  dissolve  in  the  melted  silicates  which 
form  glass,  and  it  has  been  sometimes  used  for  the  manufacture 
of  white,  opaque  glass,  but  less  expensive  materials  are  usually 
employed. 

Stannic  Acids.  Tin  resembles  silicon  in  that  several  acids 
are  derived  from  the  same  anhydride,  SnO2,  just  as  there  is  a 
long  list  of  silicic  acids  derived  .from  silicon  dioxide,  SiO2. 
Berzelius  discovered  in  1817  that  the  stannic  acid,  obtained  by 
precipitation  from  a  solution  of  stannic  chloride,  SnCU,  is  very 
different  in  its  properties  from  metastannic  acid,  which  is  formed 
by  treating  tin  with  nitric  acid,  and  his  study  of  these  compounds 
led  him  to  propose  the  word  isomer  to  designate  a  compound 
having  the  same  composition  as  some  other  compound  which 
has  different  properties.  At  that  time  the  compounds  which 
we  now  call  anhydrides  were  called  acids,  and  the  compounds 
which  he  considered  isomeric  were  the  anhydrides  of  the  two 
acids  rather  than  the  compounds  which  we  should  now  call 
stannic  and  metastannic  acids.  Later  investigations  have 
shown  that  neither  the  free  acids,  when  dried  in  the  air,  nor 
their  salts  are  isomeric  as  the  term  is  used  to-day.  This  will  be 
clear  from  the  following  table.  A  third  acid,  parastannic  acid, 
which  was  also  discovered  by  Berzelius,  is  included.  Stannic 
and  metastannic  acid  are  isomeric  when  dried  in  a  vacuum. 


NAME 

FORMULA  DRIED 

IN  THE  AlR 

FORMULA  DRIED 
IN  A  VACUUM 

FORMULA  OF 
POTASSIUM  SALT 

CHLORIDE 
FORMED  WITH 
HC1 

Stannic  acid 

H2SnOs.H2O 

EhSnOs 

K^nOa-EWD 

SnCU 

Metastannic 

acid  .    .    . 

HzSnsOn.Q  H2O 

H2Sn5On~4  H2O 

K2Sn6On.4  H2O 

SmOoCh  .4H2O 

Parastannic 

acid  .    .    . 

HzSnsOn.?  HjO 

H2SnsOii.2  H2O 

K2Sn6On.2or3H2O 

SniiO»CI2.2H2O 

Stannic  Acid,  H2SnO3.H2O,  or  H4SnO4,  is  obtained  by  pre- 
cipitating a  solution  of  stannic  chloride,  SnCU,  with  ammonia 
or  with  calcium  carbonate.  It  dissolves  easily  in  strong  acids 
or  in  alkalies.  From  its  solution  in  alkalies  it  is  reprecipitated 


512  A  TEXTBOOK  OF  CHEMISTRY 

i 

by  acids.  On  drying  it  is  partly  changed  to  metastannic  acid, 
and  a  failure  to  understand  this  has  led  to  much  confusion  in 
the  literature. 

Metastannic  Acid,  H2Sn5On.9  H2O  or  (H4SnO4)5,  is  the  prin- 
cipal product  formed  by  the  action  of  warm  nitric  acid  on  tin. 
If  the  mixture  obtained  in  this  way  is  dissolved  in  a  little  sodium 
hydroxide,  the  addition  of  an  excess  of  the  alkali  will  cause  the 
precipitation  of  sodium  metastannate,  while  the  sodium  stannate 
will  remain  in  solution. 

Metastannic  acid  is  insoluble  in  nitric  acid  or  sulfuric  acid. 
When  treated  with  concentrated  hydrochloric  acid  it  forms  a 
chloride,  Sn5O9Cl2.4  H2O,  which  dissolves  in  water  but  is  re- 
precipitated  by  concentrated  hydrochloric  acid.  Solutions  of 
stannic  chloride,  SnCU,  which  have  stood  for  some  time,  con- 
tain this  compound,  formed  by  hydrolysis  and  rearrangement 
or  condensation.  It  is  properly  named  metastannyl  chloride. 

*  Parastannic  Acid,  H2Sn5Ou.7  H2O,  was  obtained  by  Ber- 
zelius  by  heating  metastannic  acid  with  water  at  100°.  It  is 
quite  similar  to  metastannic  acid. 

Stannic  Chloride,  SnCl4,  is  a  volatile  liquid  which  boils  at 
114°  and  fumes  strongly  in  the  air,  owing  to  its  hydrolysis  by 
the  moisture  of  the  air  and  the  escape  of  hydrochloric  acid.  It 
dissolves  in  water  and  forms  several  hydrates,  but  it  seems  pretty 
certain -that  these  contain  compounds  of  the  same  general  char- 
acter as  metastannyl  chloride  (see  above)  rather  than  hydrates 
of  stannic  chloride,  as  that  would  ordinarily  be  understood. 
But  our  knowledge  of  the  structure  of  hydrates  in  general  is 
still  very  imperfect. 

Stannic  Sulfide,  SnS2,  separates  as  a  yellow,  amorphous  pre- 
cipitate when  hydrogen  sulfide  is  passed  into  an  acid  solution 
of  a  stannic  salt.  It  dissolves  in  concentrated  hydrochloric 
acid,  resembling  antimony  and  differing  from  arsenic  in  this 
respect.  It  also  is  not  precipitated  from  and  dissolves  in  solu- 
tions of  oxalic  acid  from  which  antimony  can  be  precipitated  as 
the  sulfide,  Sb2Ss.  Stannic  sulfide  dissolves  in  ammonium  sulfide 
as  ammonium  sulfostannate,  (NH4)2SnS3. 


FIREPROOFING  COTTON  GOODS  513 

Fireproofing  of  Cotton  Goods.  Many  fatal  accidents  occur 
every  year  from  the  burning  of  clothing,  and  serious  accidents 
have  occurred  from  the  burning  of  curtains  and  fabrics  in  theaters. 
A  variety  of  substances  have  been  used  to  render  fabrics  less 
inflammable.  One  of  the  best  of  these  is  stannic  oxide.  Pro- 
fessor Perkin  describes  its  application  as  follows :  "  The  flan- 
nelette (or  other  material)  is  run  through  a  solution  of  sodium 
stannate  of  approximately  45°  Tw  (sp.  gr.  1.225)  in  such  a  man- 
ner that  it  becomes  thoroughly  impregnated.  It  is  then 
squeezed  to  remove  the  excess  of  stannate  solution,  passed  over 
heated  copper  drums  to  thoroughly  dry  it,  after  which  it  is 
run  through  a  solution  of  ammonium  sulfate  of  about  15°  Tw, 
and  again  squeezed  and  dried.  Apart  from  the  precipitated 
stannic  oxide,  the  material  now  contains  sodium  sulfate  and  this 
is  removed  by  passage  through  water;  the  material  is  then 
dried  and  subjected  to  the  ordinary  processes  of  finishing.  A 
long  series  of  trials,  carried  out  under  the  most  stringent  condi- 
tions, have  conclusively  proved  that  material,  subjected  to  this 
process,  is  permanently  fireproof.  No  amount  of  washing 
with  hot  soap  and  water  will  remove  the  fireproofing  agent, 
or  in  other  words,  the  property  of  resisting  flame  lasts  so 
long  as  the  material  itself  lasts."  (Address  before  the  Inter- 
national Congress  of  Applied  Chemistry,  New  York  City,  Sep- 
tember, 1912.) 

Lead,  Pb,  207.1.  Occurrence,  Metallurgy.  Lead  is  most 
often  found  as  the  sulfide,  PbS,  in  the  form  of  galena,  a  heavy 
black  mineral  which  crystallizes  in  cubes  having  a  bright,  metallic 
luster.  Galena  is  usually  associated  with  other  minerals,  es- 
pecially with  sphalerite,  ZnS,  and  pyrites,  FeS2,  and  with  a 
gangue  of  quartz,  SiC>2,  fluorite,  CaF2,  barite,  BaSCh,  or  calcite, 
CaCO3. 

Lead  is  obtained  from  the  ore  by  roasting  it  to  convert  a  part 
of  the  sulfide  to  the  oxide  or  sulfate : 

PbS  +  3  O  =  PbO  +  SO2 
PbS  +  4  O  =  PbS04 


514  A  TEXTBOOK  OF  CHEMISTRY 

If  the  mixture  of  sulfide  with  oxide  or  sulfate  is  heated  with 
exclusion  of  air  or  in  a  reducing  atmosphere,  the  compounds 
mutually  reduce  each  other  and  sulfur  dioxide  escapes : 
2  PbO  +  PbS  =  3  Pb  +  SO2 
PbS04  +  PbS  =  2  Pb  +  2  SO2 

This  process  is  most  suitable  for  very  pure  ores.  Less  pure  ores, 
which  are  often  reduced  for  the  silver  and  other  metals  which 
they  contain  rather  than  for  the  lead,  are  usually  reduced  in  a 
blast  furnace  (p.  541)  by  the  combined  action  of  coal  or  coke 
and  iron.  The  iron  combines  with  the  sulfur  of  the  galena, 
reducing  it  to  metallic  lead : 

PbS  +  Fe  =  FeS  +  Pb 

The  recovery  of  silver,  which  is  usually  present  in  crude  lead, 
has  been  discussed  in  a  previous  chapter. 

Properties  and  Uses  of  Lead.  Alloys.  Lead  is  the  heaviest 
of  the  cheaper  metals,  having  a  specific  gravity  of  1 1 .34,  which 
is  even  higher  than  that  of  silver.  It  is  because  of  this  property, 
and  also  because  of  the  ease  with  which  lead  can  be  melted  and 
cast,  that  it  is  used  for  bullets  and  shot.  Lead  melts  at  327.4°. 
It  has  a  bright  white  luster  when  freshly  cut,  but  tarnishes 
quickly  in  the  air.  Lead  dissolves  easily  in  nitric  acid  but  only 
slowly  and  to  a  slight  extent  in  hydrochloric  acid.  Even  hot 
sulfuric  acid  scarcely  attacks  the  metal  till  it  has  the  specific 
gravity  of  1.72.  A  more  concentrated  acid  dissolves  lead  sulfate 
and  attacks  metallic  lead  strongly. 

Lead  is  so  soft  that  it  can  be  pressed  through  a  die  into  the 
form  of  tubing,  by  means  of  hydraulic  pressure.  Such  tubing 
is  used  for  waste  pipes  from  sinks  and  for  similar  purposes  and 
is  liked  by  plumbers  because  of  the  ease  with  which  it  can  be 
worked.  It  is  not  suitable  for  pipes  to  convey  drinking  water, 
because  a  little  of  the  lead  is  liable  to  dissolve,  and  all  soluble 
lead  compounds  are  very  poisonous.  Lead  has  the  further, 
very  dangerous,  property,  that  it  acts  as  a  cumulative  poison  so 
that  minute  quantities  taken  daily  for  some  weeks  or  months 
may  finally  produce  fatal  results. 


LEAD  515 

Lead  has  such  slight  tenacity  that  it  cannot  be  drawn  into 
wire,  in  spite  of  its  softness.  It  can  be  rolled  into  sheets  and 
beaten  into  thin  foil.  Foil  made  from  the  alloy  with  tin  has 
sometimes  been  substituted  for  pure  tin  foil  to  wrap  around 
articles  of  food,  but  such  use  is  strongly  condemned  because  of 
the  poisonous  character  of  the  lead. 

The  more  important  alloys  of  lead,  solder,  Britannia  metal, 
pewter,  Babbitt  and  other  antifriction  metals,  type  metal, 
stereotype  metal  and  fusible  alloys  used  for  safety  fuses,  have 
been  mentioned  in  previous  chapters. 

Oxides  of  Lead.  There  are  three  definite,  well-characterized 
oxides  of  lead :  lead  monoxide,  or  litharge,  PbO,  lead  plumbate, 
usually  called  red  lead  or  minium,  Pb3O4,  and  lead  dioxide,1  or 
plumbic  anhydride,  PbO2.  Three  other  oxides,  Pb2O,  PbsOy  and 
Pb2Oa,  have  been  described  by  various  authors,  but  there  is 
considerable  doubt  whether  these  represent  definite  compounds 
or  not.  The  evidence  in  favor  of  the  existence  of  the  oxide, 
Pb2O3  (or  PbPbO3,  lead  metaplumbate),  is  better  than  that  for 
the  other  two. 

Lead  Monoxide  or  Litharge,  PbO,  is  readily  formed  by  exposing 
lead  at  a  red  heat  to  the  action  of  the  air.  At  that  tempera- 
ture the  melted  film  of  litharge  constantly  flows  to  the  side,  ex- 
posing a  fresh  surface  to  oxidation.  The  melted  litharge  is 
collected,  allowed  to  solidify,  and  is  then  ground  to  a  fine  powder 
for  the  market.  It  forms  a  buff-colored  powder  used  in  the  manu- 
facture of  "  boiled  "  linseed  oil  and  flint  glass,  and  for  other 
purposes. 

For  the  manufacture  of  red  lead,  metallic  lead  is  first  oxidized 
at  a  lower  temperature,  such  that  the  oxide  formed  does  not 
melt.  This  oxide  is  then  heated  to  dull  redness,  but  below  the 
melting  point  of  litharge,  with  free  access  of  air. 

The  chemical  character  of  red  lead  is  most  clearly  shown  by 
the  action  of  dilute  nitric  acid  upon  it.  This  dissolves  two 

1  Often  called  lead  peroxide.  It  is  better  to  restrict  the  designa- 
tion peroxide  to  compounds  having  a  structure  similar  to  that  of 
hydrogen  peroxide,  H — O — O — H. 


516  A  TEXTBOOK  OF  CHEMISTRY 

thirds  of  the  lead  and  leaves  a  dark  brown  residue  of  lead  di- 
oxide, or  plumbic  anhydride,  PbO2  : 

Pb3O4  +  4  HNO3  =  2  Pb(NO3)2  +  PbO2  +  2  H2O 

This  reaction  shows  that  two  thirds  of  the  lead  is  basic  and 
one  third  acidic  in  character,  or,  in  other  words,  that  red  lead 
is  a  lead  salt  of  plumbic  acid,  H4PbO4.  The  relation  is  clearer 
if  we  write  the  formula  Pb2PbO4.  When  this  is  treated  with 
nitric  acid,  the  two  lead  atoms  which  form  the  positive  ions, 
Pb++,  of  the  salt,  are  easily  exchanged  for  the  hydrogen  of  the 
acid,  but  the  plumbic  acid,  H4PbO4,  which  results,  is  unstable 
and  decomposes  at  once  to  water  and  plumbic  anhydride  or 
lead  dioxide,  PbO2,  just  as  carbonic  acid,  H2CO3,  decomposes 
to  carbon  dioxide  and  water. 

Red  lead  is  used  as  a  pigment,  as  an  oxidizing  agent  in  glass 
manufacture  and  with  linseed  oil  as  a  lute  in  plumbing. 

As  litharge  is  oxidized  to  lead  plumbate,  Pb2PbO4,  a  mixture  of 
lime,  CaO,  and  litharge  may  be  readily  oxidized  at  low  redness 
to  calcium  orthoplumbate,  Ca2PbO4.  Dilute  acids  decompose 
this  with  the  separation  of  lead  peroxide. 

Storage  Batteries.  A  storage  battery  which  has  been  charged 
contains  two  kinds  of  lead,  one  of  which  has  been  more  or  less 
completely  changed  to  lead  dioxide,  while  the  other  consists 
of  spongy,  metallic  lead.  During  the  discharge  the  electrons 
escape  to  the  connecting  wire  from  the  metallic  lead,  Pb,  chang- 
ing it  to  lead  ions,  Pb++,  which  combine  with  sulfate  ions  of  the 
solution,  forming  insoluble  lead  sulfate,  PbSO4 : 

Pb  +  SO4"  =  PbSO4  +  2- 

The  electrons  pass  through  the  plate  and  connections  to  the 
other  plate.  At  the  other  plate  the  two  electrons  combine  with 
the  lead  of  the  lead  dioxide,  reducing  it  from  the  quadrivalent 
to  the  bivalent  form,  resulting  in  the  formation  of  lead  sulfate 
and  the  liberation  of  a  sulfate  ion,  SO4~~,  from  the  sulfuric 
acid: 

2-  +  PbO2  +  2  H2SO4  =  PbSO4  +  2  H2O  +  SO4~ 


STORAGE  BATTERIES  517 

or  Pb++++O--O---  +  2-  =  Pb++O--  -f-  O" 

O—  +  H2SO4  =  H20  +  SO4= 
and  Pb++O-"  +  H2S04  =  Pb++SO4=  +  H2O 

During  the  discharge  there  is  a  difference  of  potential  of 
about  two  volts  between  the  two  plates,  and  the  passage  of  the 
current  develops  a  very  considerable  amount  of  electrical 
energy.  In  charging  the  battery  the  reverse  operations  take 
place.  At  the  cathode,  which  is  connected  with  the  negative 
pole  of  the  dynamo,  electrons  combine  with  lead  ions,  Pb+4", 
reducing  them  to  metallic  lead,  Pb,  and  leaving  tWe  sulfate  ions, 
SO4~~,  of  the  lead  sulfate  free  to  pass  into  solution. 

Pb++SO4—  +  2-  =  Pb  +  SO4~- 

At  the  anode  the  bivalent  lead  ions,  Pb++,  lose  two  electrons, 
giving  tetravalent  lead  ions,  Pb++++,  which  combine  momentarily 
with  another  sulfate  ion,  SO4~~~,  to  form  lead  tetrasulfate, 
Pb(SO4)2.  The  lead  tetrasulfate  is  at  once  hydrolyzed  to  lead 
peroxide,  PbO2,  and  sulfuric  acid. 

Pb++SO4-"  +  SO4~-  =  Pb++++(SO4--)2  +  2- 
Pb(SO4)2  +  2  H2O  =  Pb02  +  2  H2SO4 

It  will  be  noticed  that  in  both  processes  sulfate  ions,  SO4=, 
are  discharged  at  one  plate  and  enter  into  combination  with 
the  lead,  while  at  the  other  plate  sulfate  ions  pass  into  solution. 
In  charging  the  battery  the  discharged  sulfate  ions  combine 
with  the  lead  of  lead  sulfate,  PbSO4,  forming  the  tetrasulfate, 
Pb(SO4)2.  In  discharging  the  discharged  sulfate  ions  combine 
with  the  metallic  lead,  forming  lead  sulfate.  At  the  same  time 
sulfate  ions  must,  of  course,  migrate  through  the  solution  in  one 
direction  in  charging,  in  the  opposite  direction  in  discharging. 

The  charging  is,  of  course,  accompanied  by  an  absorp- 
tion of  energy.  Practically,  a  very  high  efficiency  can  be  se- 
cured, the  energy  obtained  during  the  discharge  approaching 
closely  to  that  absorbed  in  charging.  Since  the  sulfate  radicals 
are  mostly  in  the  form  of  sulfuric  acid  in  the  charged  battery 


518  A  TEXTBOOK  OF  CHEMISTRY 

and  in  the  form  of  lead  sulfate  in  the  battery  after  discharge, 
the  amount  of  sulfuric  acid  in  the  electrolyte,  which  can  be  easily 
determined  by  a  hydrometer,  furnishes  a  pretty  close  indication 
of  the  condition  of  the  cell. 

Lead  dioxide  when  warmed  with  hydrochloric  acid  gives  at 
first  lead  tetrachloride,  PbCU,  but  this  is  unstable  and  decom- 
poses into  lead  chloride,  PbCl2,  and  chlorine,  C12,  exactly  as 
manganese  tetrachloride  does  (p.  101).  It  is  noteworthy  that 
this  is  entirely  different  from  the  conduct  of  barium  peroxide, 
BaO2,  or  sodium  peroxide,  Na2O2,  either  of  which  gives  hydrogen 
peroxide,  H2t)2,  with  hydrochloric  acid.  (What  does  this  dif- 
ference in  conduct  indicate  as  to  the  structure  of  these  three 
oxides?) 

Lead  Sulfide,  PbS,  is  found  in  nature  as  the  mineral  galena 
and  is  formed  as  a  black  precipitate  by  the  action  of  hydrogen 
sulfide  on  a  solution  of  a  lead  salt  in  a  dilute  acid.  It  is  not 
precipitated  in  the  presence  of  much  hydrochloric  acid,  or  of 
much  of  any  other  strong  acid.  It  dissolves  easily  in  nitric 
acid. 

Lead  Chloride,  PbCl2,  forms  as  a  white,  crystalline  precipi- 
tate on  adding  hydrochloric  acid  to  a  solution  of  almost  any 
soluble  salt  of  lead.  It  dissolves  in  125  parts  of  water  at  18° 
and  in  30  parts  of  boiling  water.  It  is  less  soluble  in  dilute 
hydrochloric  acid  than  in  pure  water,  but  is  more  soluble  in 
concentrated  acid,  doubtless  because  of  the  formation  of  a 
complex  compound  with  the  acid,  such  as  chloroplumbous  acid, 
H2PbCl4. 

Lead  Tetrachloride,  PbCU.  By  dissolving  lead  peroxide, 
PbO2,  in  cold,  concentrated  hydrochloric  acid  a  solution  con- 
taining lead  tetrachloride,  or  more  likely  chloroplumbic  acid, 
H2PbCle,  is  obtained.  On  the  addition  of  ammonium  chloride, 
ammonium  chloroplumbate,  (NH4)2PbCl6,  separates.  If  this 
salt  is  dissolved  in  cold  concentrated  sulfuric  acid,  hydrochloric 
acid  escapes  and  lead  tetrachloride  separates  below  the  acid  as  a 
heavy  yellow  liquid.  It  is  quite  unstable,  decomposing  readily 
into  lead  chloride,  PbCl2,  and  chlorine.  It  is  hydrolyzed  by 


LEAD  SALTS  519 

water  to  lead  dioxide  and  hydrochloric  acid.  With  a  little 
hydrochloric  acid  it  gives  a  yellow  compound,  chloroplumbic 
acid,  H2PbCl6.  A  considerable  number  of  salts  of  this  acid  are 
known.  The  ammonium  salt  is  mentioned  above.  These  salts 
are  usually  called  double  salts  of  lead  tetrachloride  with  other 
chlorides,  and  are  frequently  written  in  the  form  PbCU.2  KC1 
instead  of  K2PbCl6. 

Lead  Sulfate,  PbSO4,  may  be  prepared  by  the  precipitation 
of  any  soluble  salt  of  lead  with  dilute  sulfuric  acid.  It  may 
also  be  obtained  by  roasting  lead  sulfide  at  a  moderate  tem- 
perature. The  compound  prepared  in  the  latter  manner  has 
been  used  in  America  to  a  limited  extent  as  a  pigment  in  place 
of  the  ordinary  white  lead  described  below.  It  is  a  difficultly 
soluble  salt,  but  is  distinctly  more  soluble  than  barium  sulfate. 
It  dissolves  easily  in  a  solution  of  ammonium  acetate. 

Lead  Nitrate,  Pb(NO3)2,  is  a  very  easily  soluble  salt  which 
may  be  prepared  by  dissolving  either  lead  or  litharge  in  dilute 
nitric  acid.  Either  lead  nitrate  or  lead  acetate  may  be  used 
for  the  preparation  of  the  insoluble  salts  of  lead,  especially  of 
the  chromate,  PbCrO4. 

Lead  Acetate  or  Sugar  of  Lead,  Pb(C2H3O2)2.3  H2O,  is  an 
easily  soluble  salt  prepared  by  dissolving  litharge  in  acetic 
acid.  Sugar  of  lead  has  sometimes  been  used  in  hair  dyes,  but 
its  use  in  this  way  is  considered  dangerous  and  liable  to  cause 
paralysis. 

Basic  Lead  Acetates,  Pb(C2H3O2)OH  and  Pb(C2H3O2)2. 
2  Pb(OH)2,  are  formed  by  dissolving  litharge,  PbO,  in  a  solu- 
tion of  lead  acetate.  Such  a  solution  is  used  to  clarify  dark- 
colored  sugar  solutions  to  prepare  them  for  determinations  with 
the  polarimeter. 

Lead  Carbonate,  PbCO3.  The  normal  salt  may  be  precipi- 
tated by  adding  sodium  carbonate  to  a  solution  of  lead  acetate. 
It  is  so  difficultly  soluble  that  it  can  also  be  precipitated  by 
passing  carbon  dioxide  through  a  solution  of  lead  acetate,  and 
lead  acetate  gives  a  turbid  solution  with  ordinary  distilled  water 
because  of  its  formation. 


520  A  TEXTBOOK  OF  CHEMISTRY 

Basic  Lead  Carbonate,  or  White  Lead,  Pb3(CO3)2(OH)2  or 
2  PbCO3.Pb(OH)2,  has  been  manufactured  for  many  centuries 
for  use  as  a  pigment.  The  principal  process  used,  known  as  the 
"  Dutch  process,"  has  been  scarcely  changed  in  principle  for  a 
very  long  time. 

Plates,  or  "  buckles,"  of  lead  about  one  eighth  of  an  inch 
thick  and  five  and  one  half  inches  in  diameter  are  cast  in  the 
form  shown  in  Fig.  104.  These  plates  are  packed  in  pots,  Fig. 
105,  having  about  250  cc.  of  dilute  acetic  acid  or  vinegar  in  the 


Fig.  104  Fig.  105 

t 

bottom.  These  pots  are  then  packed  in  layers  with  alternate 
layers  of  spent  tanbark  until  a  large  room  is  filled.  The  room 
is  left  to  itself  for  about  three  months.  The  combined  action 
of  the  vapors  of  acetic  acid  and  air  on  the  lead  plates  causes 
them  to  corrode  superficially,  and  the  basic  acetate  formed  is 
changed  to  carbonate  by  the  carbon  dioxide  which  comes  from 
the  fermentation  of  the  tanbark.  As  the  carbonate  is  formed 
some  acetic  acid  is  continuously  liberated,  and  this,  with  the 
air,  carries  on  the  corrosion  till  the  "  buckles  "  are  almost 
completely  changed  to  basic  carbonate.  At  the  end  of  three 
months  the  pots  are  emptied,  the  white  lead  is  finely  ground 
and  separated  from  particles  of  lead  an4  coarse  particles  of 
material  by  bolting  and  lixiviating  with  water.  The  fine  powder 
is  then  dried  and  intimately  incorporated  with  linseed  oil  for 


WHITE   LEAD  521 

the  market.  Another  process  is  to  mix  the  moist  powder 
directly  with  linseed  oil.  Owing  to  the  relation  between  the 
surface  tension  of  white  lead  toward  water  and  that  toward 
linseed  oil,  the  latter  is  able  to  displace  the  water  and  joins  with 
the  white  lead  to  form  a  paint  which  is  practically  identical  in 
composition  and  properties  with  the  pigment  prepared  from  the 
dry  powder.  Owing  to  the  poisonous  character  of  lead  com- 
pounds, workmen  in  white  lead  factories  and  painters  often 
suffer  from  a  painful  and  sometimes  fatal  disease,  called  lead 
colic.  Stringent  laws  have  been  passed  by  some  states  for  the 
protection  of  the  workmen  from  inhaling  the  dust  and  from 
poisoning  in  other  ways.  Some  manufacturers  spend  large  sums 
of  money  to  protect  their  workmen  from  the  danger  of  poisoning. 
White  lead  depends  for  its  value  on  the  fact  that  it  is  an 
amorphous,  very  fine,  opaque  white  powder,  the  opacity  being 
much  greater  than  that  of  barium  sulfate,  which  is  sometimes 
used  as  an  adulterant  or  substitute.  The  fact  that  water  will 
not  wet  it  when  it  is  in  contact  with  linseed  oil  is  also  a  factor 
of  prime  importance.  White  lead  is  blackened  by  hydrogen 
sulfide  and  for  that  reason  is  less  suitable  than  zinc  white  and 
lithopone  for  use  in  chemical  laboratories  or  in  localities  where 
it  is  subjected  to  the  action  of  sewer  gas. 


CHAPTER  XXX 
VANADIUM   AND    CHROMIUM   GROUPS 

Group  V.  Vanadium,  columbium  (or  niobium)  and  tantalum 
are  the  elements  of  Group  V,  which  alternate  with  phosphorus, 
arsenic,  antimony  and  bismuth.  They  are  usually  classed  as 
rare  elements,  but  vanadium  and  tantalum  have  acquired  some 
industrial  importance.  Each  forms  a  pentoxide,  corresponding 
to  P2O5,  and  salts  of  acids  corresponding  to  themetaphosphates, 
MPO3,  pyrophosphates,  M4P2O7,  and  orthophosphates,  M3PO4. 
Salts  of  the  form  MgV^C^  are  also  known  for  each. 

Vanadium,  V,  51.0,  is  very  widely  diffused  in  nature,  being 
found  in  small  amounts  in  almost  all  clays  and  massive  rocks. 
The  most  important  mineral  is  vanadinite,  Pb5(VO4)3Cl,  which 
corresponds  in  composition  to  apatite,  Ca5(PO4)3F. 

Vanadium  is  a  silvery  white  metal  having  a  specific  gravity 
of  5.5  and  melting  at  about  1720°.  It  is  used  as  an  addition  to 
steel,  increasing  its  hardness,  malleability  and  tensile  strength. 
Ferrovanadium  is  an  alloy  with  iron,  which  is  much  more  easily 
prepared  than  the  pure  metal.  It  is  the  commercial  form  used 
for  addition  to  steel. 

*  Vanadium  forms  compounds  in  which  it  is  bivalent,  tri- 
valent,  quadrivalent  and  quinquivalent.  The  following  may 
be  mentioned :  vanadous  chloride,  VC12,  vanadous  sulfate, 
VSO4.7  H2O,  vanadous  sulfide,  VS,  vanadic  chloride,  VC13, 
vanadic  sulfide,  V2S3,  vanadic  sulfate,  V2(SO4)3,  vanadium  alum, 
KV(SO4)2.12  H2O,  vanadium  tetrachloride,  VC14,  vanadium 
pentachloride,  VCls,  vanadium  oxychloride,  VOC13,  sodium 
orthovanadate,  Na3VO4.12H2O,  sodium  pyrovanadate, 
Na2V2O7.18H2O,  sodium  metavanadate,  NaVO3.2H2O.  More 
complex  vanadates  and  other  complex  compounds  of  a  great 
variety  of  forms  have  been  prepared. 

522 


COLUMBIUM.    TANTALUM  523 

*  Columbium,  Cb,93.5  (or  Niobium,  Nb).     In  1801  Hatchett 
discovered  a  new  element  in  a  mineral  from  Haddam,  Con- 
necticut.    He  called  the  mineral   columbite  (from    Columbia, 
the  poetical  name  for  America),  and  the  element  columbium. 
It  is  probable  that  the  compounds  which  he  prepared  contained 
both  columbium  and  tantalum,  and  the  two  elements  were  first 
clearly  separated  and  characterized  by  H.  Rose  in  1844.     Rose 
either  overlooked  or  ignored  the  discovery  of  Hatchett  and 
called  the  two  metals  tantalum  and  niobium. 

Columbite,  the  mineral  in  which  columbium  was  discovered, 
is  a  ferrous  metacolumbate,  Fe(CbO3)2,  containing  ferrous 
metatantalate,  Fe(TaO3)2,  in  isomorphous  mixture.  The 
formula  is  more  properly  written,  Fe((Cb,Ta)O3)  or,  in 
the  oxide  form,  which  is  most  frequently  used  by  mineralogists, 
FeO((Cb2,Ta2)05). 

Elementary  columbium  is  still  more  metallic  in  its  properties 
than  vanadium.  It  has  a  specific  gravity  of  7.4  and  melts  at 
about  2200°.  The  following  are  typical  compounds :  colum- 
bium trichloride,  CbCl3,  columbium  pentafluoride,  CbFs,  colum- 
bium oxyfluoride,  CbOF3,  columbium  pentoxide,  Cb2O5,  mag- 
nesium orthocolumbate,  Mg3(CbC>4)2,  calcium  pyrocolumbate, 
Ca2Cb2O7,  potassium  hexacolumbate,  4  K2O.3  Cb2O5.16  H2O  or 
K8Cb2Og.l6  H2O.  Many  complex  columbates  ana  other  com- 
plex compounds  are  known. 

*  Tantalum,  Ta,  181.5.     In  1902  Ekeberg  examined  a  mineral 
from  Finland  which   closely  resembles  columbite  and  gave  to 
it  the  name  tantalite.     He  called  the  element  which  it  contains 
tantalum,  but  the  compounds  which  he  prepared  were  doubtless 
mixtures    containing    both    columbium    and    tantalum.     The 
mineral  which  he  studied  always  contains  both  elements,  and 
it  is  most  properly  called  columbite  when  the  columbium  is  in 
excess  and  tantalite  when  there  is  more  of  the  tantalum.     As 
stated  above,  H.  Rose  first  distinguished  sharply  between  the 
two  elements  and  cleared  up  the  confusion  of  the  earlier  workers. 

Tantalum  is  a  bright  metal  somewhat  resembling  platinum 
in  appearance.  It  is  ductile  and  can  be  drawn  into  fine  wire. 


524  A  TEXTBOOK  OF  CHEMISTRY 

As  the  melting  point  is  2850°  it  is  suitable  for  the  filaments  of 
electric  lights,  and  was  used  for  a  short  time  in  that  way,  but 
was  quickly  displaced  by  tungsten.  The  specific  gravity  of  the 
metal  is  16.5. 

Among  the  compounds  of  tantalum  are  the  potassium  fluo- 
tantalate,  2KF.TaF5  or  K2TaF7,  by  means  of  which  tantalum 
can  be  best  separated  from  columbium  (E.  F.  Smith),  tantalum 
pentachloride,  Tads,  tantalum  pentoxide,  Ta2Os,  and  a  series 
of  tantalates. 

Group  VI.  Chromium,  molybdenum,  tungsten  and  uranium 
alternate  with  sulfur,  selenium  and  tellurium  of  Group  VI, 
exactly  as  the  three  elements  last  considered  alternate  with  the 
elements  of  the  phosphorus  family. 

The  elements  form  compounds  in  which  they  appear  bivalent, 
trivalent  and  sexivalent,  especially,  but  also  some  in  which 
they  are  quadrivalent.  They  resemble  the  sulfur  family  in 
such  compounds  as  potassium  chromate,  K2CrO4,  potassium 
dichromate,  K2Cr2O7,  and  ammonium  molybdate,  (NH4)2MoO4. 

Chromium,  Cr,  52.0,  is  found  chiefly  as  ferrous  chromite, 
FeO.C^Os  or  Fe(CrO2)2,  which  is  known  as  chromite  or  chrome 
iron  ore,  a  black  mineral  isomorphous  with  magnetite,  Fe3O4  or 
(FeO.Fe2O3).  Chromium  is  also  found  as  lead  chromate, 
PbCrO4,  and  was  first  discovered  in  that  mineral  by  Vauquelin 
in  1797.  The  name  was  given  because  of  the  colored  compounds 
which  it  forms. 

Metallurgy,  Uses.  Chromium  is  now  prepared  by  Gold- 
schmidt's  thermite  process  (p.  497)  by  igniting  a  mixture  of 
chromic  oxide,  Cr2O3,  and  aluminium. 

Metallic  chromium  is  a  white,  crystalline,  extremely  hard 
metal  —  as  hard  as  corundum.  It  has  a  specific  gravity  of 
about  7.0.  It  melts  at  1520°  and  boils  at  2200°.  It  is  used  as 
an  addition  to  steel,  making  it  extremely  hard  and  resistant  to 
the  penetration  of  projectiles  when  used  for  armor  plate. 
Such  steel  requires  very  careful  heat  treatment  to  bring  out  its 
best  properties.  Alloys  with  nickel  or  cobalt  resist  the  action 
of  acids  and  are  proving  useful  for  special  purposes. 


CHROMIUM  525 

Chromous  Chloride,  CrCl2.4  H2O.  A  light  blue  solution 
containing  chromous  chloride  is  easily  prepared  by  the  action 
of  zinc  and  hydrochloric  acid  on  a  solution  of  chromic  chloride, 
CrCl3,  or  of  chromic  anhydride,  CrO3.  Air  must  be  carefully 
excluded,  as  air  or  oxygen,  in  the  presence  of  hydrochloric  acid, 
changes  the  compound  back  to  green  chromic  chloride.  The 
solid  chromous  chloride  is  white. 

Chromic  Oxide,  Cr2O3,  is  a  green  powder  formed  by  igniting 
the  hydroxide.  A  more  or  less  pure  chromic  oxide  or  hydroxide 
is  prepared  in  a  variety  of  ways  and  used  as  a  pigment  under 
the  name  of  chrome  green  or  Guignet's  green.  The  latter  con- 
tains a  little  boric  acid.  A  green  pigment  which  is  made  by 
mixing  chrome  yellow,  PbCrO4,  and  Prussian  blue,  Fe4(FeC6N6)3, 
is  also  often  sold  and  used  under  the  name  of  "  chrome  green." 

Chromic  Hydroxide,  Cr2O3.4  H2O.  From  analogy  with  the 
formula  of  the  chloride,  the  formula  Cr(OH)3  is  often  given  for 
the  hydroxide.  The  compound  formed  by  precipitating  a 
chromic  salt  with  ammonium  hydroxide  has  the  composition 
Cr2O3.4  H2O,  however,  when  it  is  dried  in  a  vacuum.  A  com- 
pound, or  mixture,  having  the  composition  Cr(OH)3  has  been 
obtained  in  some  cases,  but  it  is  not  the  usual  composition  of  the 
hydroxide.  Chromic  hydroxide  dissolves  to  some  extent  in 
alkalies,  but  does  not  show  as  marked  acidic  properties  of 
aluminium  hydroxide.  A  number  of  chromites  have  been  pre- 
pared, however,  and  ferrous  chromite,  Fe(CrO2)2,  has  been  men- 
tioned above  as  the  most  important  ore  of  chromium. 

Chromic  Chloride,  CrCl3,  was  formerly  prepared  by  heating 
a  mixture  of  chromic  oxide,  Cr2O3,  and  carbon  in  a  current  of 
chlorine.  It  sublimes  in  beautiful  reddish  pink  or  violet  leaflets. 
These  are  almost  insoluble  in  water,  but  dissolve  easily  to  a  green 
solution  in  water  containing  a  trace  of  chromous  chloride,  CrCl2. 

Hydrates  of  Chromic  Chloride.  A  solution  containing 
chromic  chloride  is  easily  prepared  by  the  reduction  of  chromic 
anhydride  or  a  solution  of  a  chromate : 

2  CrO3  +  6  HC1  +  3  C2H5OH  =  2  CrCl3  +  3  C2H4O  +  6  H2O 

Alcohol  Aldehyde 


526  A  TEXTBOOK  OF  CHEMISTRY 

Other  reducing  agents,  such  as  sulfurous  acid,  may  be  used, 
but  alcohol  has  the  advantage  that  it  leaves  no  other  compound 
in  the  solution. 

By  various  methods  it  is  possible  to  prepare  two  isomeric 
hydrates  of  chromic  chloride,  both  of  which  have,  in  the  crystal- 
line form,  a  composition  corresponding  to  the  formula 
CrCla.G  H2O.  One  of  these  hydrates  dissolves  in  water  to  a 
green  solution,  while  the  other  gives  a  violet  solution.  Each 
can  be  changed  more  or  less  completely  into  the  other,  and 
solutions  of  chromic  chloride  usually  contain  the  two  compounds 
in  equilibrium  with  each  other.  The  green  solution  has  a  lower 
electrical  conductivity  than  the  violet.  If  silver  nitrate  is  added 
to  an  ice-cold  solution  of  the  green  hexahydrate,  only  one  third 
of  the  chlorine  is  precipitated  at  first. 

The  most  satisfactory  explanation  of  these  facts  is  given  by 
the  theory  of  Werner.  He  supposes  that  in  each  compound  six 
atoms  or  groups  are  directly  combined  with  or  arranged  about 
the  chromium  atom  to  form  a  characteristic  group.  In  the 
violet  chloride  he  supposes  that  these  six  groups  are  all  mole- 
cules of  water  and  he  calls  this  hydrate  hexaaquochromic  chloride 
and  writes  the  formula  (Cr(OH2)e)Cl3.  In  the,  green  chloride, 
on  the  other  hand,  he  considers  that  four  molecules  of  water  and 
two  atoms  of  chlorine  form  the  six  groups  directly  combined 
with  the  chromium,  while  one  atom  of  chlorine  and  two  mole- 
cules of  water  are  less  directly  connected.  He  writes  the  formula 


of   the   green  chloride,  accordingly,    cr*         Cl  +  2  H2O. 

[_        (OH2J4J 

Three  chlorine  atoms  of  the  violet  chloride  are  readily 
ionized,  giving  electrical  conductivity  and  easy  precipita- 
tion as  silver  chloride.  But  only  one  of  the  chlorine  atoms 
of  the  green  chloride  ionizes  to  give  electrical  conductivity, 
and  only  this  one  atom  is  directly  precipitated  by  silver 
nitrate. 

Werner  and  others  have  shown  that  a  large  number  of  com- 
plex inorganic  compounds,  containing  water,  ammonia  and 
other  groups,  exhibit  isomerism  similar  to  that  of  these  hydrates. 


CHROMATES  527 

(See  Werner,  Neuere  Anschauung  auf  dem  Gebiet  der  anorgan- 
ischen  Chemie.) 

Potassium  Chromium  Sulfate  or  "  Chrome  Alum," 
KCr(SO4)2.12  H2O.  If  potassium  dichromate,  K2Cr2O7,  is 
warmed  with  alcohol  and  dilute  sulfuric  acid,  the  chromium  is 
reduced  to  the  form  of  chromic  sulfate,  02(804)3,  and  on  evapo- 
ration of  the  solution  and  cooling,  the  chromic  sulfate  and  potas- 
sium sulfate  will  combine  to  form  chrome  alum,  a  dark  violet 
salt,  which  is  isomorphous  with  ordinary  alum. 

Potassium  Chromate,  K2CrO4.  When  chrome  iron  ore, 
FeCr2O4,  is  roasted  with  potassium  carbonate,  the  chromium  is 
oxidized  and  forms  potassium  chromate.  This  is  in  accordance 
with  the  practically  universal  rule  that  metals  assume  a  higher 
state  of  oxidation  in  the  presence  of  a  base  than  in  the  presence 
of  an  acid : 

4  FeCr2O4  +  8  K2CO3  +  7  O2  =  2  Fe2O3  +  8  K2CrO4  +  8  CO2 

The  potassium  chromate  is  easily  soluble  and  can  be  easily 
separated  from  the  insoluble  ferric  oxide.  It  is  an  easily  soluble 
yellow  salt,  which  gives  a  lemon-yellow  solution. 

Potassium  Dichromate,  or  Pyrochromate,  K2Cr2O7,  is  formed 
on  adding  an  acid,  even  a  weak  acid,  to  a  solution  of  potassium 
chromate.  It  crystallizes  in  orange-red  crystals,  which  are 
much  less  soluble  than  the  crystals  of  the  chromate.  Potassium 
dichromate  is  the  practical  source  of  almost  all  other  chromium 
compounds  and  exceeds  all  of  the  others  in  commercial  impor- 
tance. It  is  used  as  an  oxidizing  agent,  as  a  mordant  in  dyeing, 
and  for  the  preparation  of  leather  in  chrome  tanning,  a  process 
rapidly  increasing  in  importance. 

Potassium  dichromate  corresponds  to  potassium  pyrosulfate, 
K^Oy.  The  chromium  in  it  is  in  the  same  state  of  oxidation 
as  in  the  chromate. 

Lead  Chromate,  or  Chrome  Yellow,  PbCrO4,  is  an  insoluble  salt 
obtained  by  precipitating  a  solution  of  lead  acetate  with  potassium 
dichromate.  It  is  a  brilliant  yellow  compound  which  is  used  as 
a  pigment  and  as  a  constituent  of  "chrome  green  "  (p.  525). 


528  A  TEXTBOOK  OF  CHEMISTRY 

*  Barium  Chromate,  BaCrO4,  is  easily  prepared  by  precipita- 
tion. It  is  insoluble  in  water  or  acetic  acid  but  dissolves  in  hy- 
drochloric or  nitric  acid. 

Chromium  Trioxide  or  Chromic  Anhydride,  CrOa,  separates 
in  the  form  of  dark  red  needles  on  adding  an  excess  of  concen- 
trated sulfuric  acid  to  a  saturated  solution  of  potassium 
dichromate.  When  mixed  with  sulfuric  or  other  acids,  it  is  a 
powerful  oxidizing  agent  and  is  much  used  for  that  purpose 
in  the  laboratory.  To  illustrate ;  such  a  mixture  with  sulfuric 
acid  will  oxidize  the  carbon  or  graphite  of  cast  iron  or  steel  to 
carbon  dioxide,  and  this  method  is  often  used  for  the  determi- 
nation of  these  elements. 

Chromyl  Chloride,  CrO2Cl2,  is  prepared  by  distilling  a  mixture 
of  potassium  dichromate,  salt  and  sulfuric  acid.  It  is  a  volatile, 
fuming,  dark  red  liquid,  which  is  a  very  powerful  oxidizing 
agent.  Many  organic  compounds  react  with  it  with  explosive 

violence.     It  may  be  considered  as  chromic  acid, 

H- 

in  which  the  two  hydroxyl  groups  have  been  replaced  by  chlorine, 
Ck          ,& 

y>Cr^    .     It  is  readily  hydrolyzed  by  water  to  hydrochloric 
Cl 

y 

\ 
and  dichromic  acids.     The  later  may  be  written  /O  . 

\ 


<V 

0> 


Molybdenum,  Mo,  96,  occurs  chiefly  as  molybdenite,  MoS2, 
a  black  mineral  closely  resembling  graphite  in  appearance  but 
having  more  than  twice  the  specific  gravity  of  that  mineral. 
The  metal  is  silver-white,  after  melting,  and  has  a  crystalline 
fracture.  It  melts  at  about  2500°. 

*  Molybdium  Trioxide,  or  Molybdic  Anhydride,  MoOs,  is  a 
yellowish  white  or  white  powder,  which  dissolves  easily  in  alka- 
lies, forming  molybdates,  such  as  sodium  molybdate,  Na2MoO4. 


MOLYBDENUM  529 

It  is  nearly  insoluble  in  acids,  but  if  a  solution  of  ammonium 
molybdate,  (NH4)2MoO4,  is  poured  into  dilute  nitric  acid  (not 
the  reverse),  the  molybdic  acid  or  anhydride  which  is  formed 
remains  in  solution  either  as  a  colloid  or  in  a  supersaturated  solu- 
tion. Such  a  solution,  known  as  "  molybdic  solution/'  is  used 
to  precipitate  solutions  of  orthophosphoric  acid,  H3PO4.  On 
adding  an  excess  of  the  molybdic  solution  to  a  solution  contain- 
ing a  soluble  orthophosphate,  or  orthophosphoric  acid,  a  very 
difficultly  soluble,  yellow  precipitate  separates.  This  has  the 
composition  (NH4)3PO4.12  MoO3.nH2O.  As  the  molybdic  an- 
hydride will  neutralize  alkalies,  the  amount  of  the  molybde- 
num and  so,  indirectly,  that  of  the  phosphorus  may  be  de- 
termined by  titration  with  a  standard  solution  of  potassium 
hydroxide : 

MoO3  +  2  KOH  =  K2MoO4  +  H2O 

If  the  precipitate  is  dissolved  in  ammonia  and  an  excess  of  dilute 
sulfuric  acid  added,  the  molybdenum  maybe  reduced  to  molybdic 
sulfate,  Mo2 (804)3,  by  means  of  zinc.  The  solution  obtained  in 
this  manner  may  be  oxidized  quantitatively  to  a  solution  con- 
taining molybdic  anhydride  by  potassium  permanganate.  This, 
again,  gives  an  indirect  determination  of  the  amount  of  phos- 
phorus. The  solution  of  molybdic  sulfate  is  very  easily  oxi- 
dized by  the  air  and  the  success  of  the  operation  depends  on  rapid 
work. 

*  Compounds  of  Molybdenum.     Molybdenum  forms  a  series 
of  oxides,  the  most  important  or  best  characterized  being,  Mo2Os, 
Mo5Oi2,  Mo2O5,  Mo3O8,  MoO2  and  MoO3.     The  sulfides  are 
Mo2S3,  MoS2,  MoS3  and  MoS4. 

*  Molybdic  anhydride,  MoO3,  forms  a  bewildering  variety  of 
complex  compounds  of  which  the  ammonium  phosphomolyb- 
date  used  in  analysis  and  referred  to  above  is  an  isolated  exam- 
ple.     Among  the    related    compounds  are    many    ammonium 
molybdates,       such       as      triammonium       dodekamolybdate 
12  MoOs.3  NHs.12  H2O,  and  phosphomolybdic  acids,  as  phos- 
phoduodecimolybdic  acid,  24  MoO3.P2O5.4  H2O. 


530  A  TEXTBOOK  OF  CHEMISTRY 

A  number  of  different  ammonium  phosphomolybdates  have 
also  been  prepared;  and  it  is  only  by  careful  attention  to  the 
proper  conditions  of  temperature,  acidity  and  concentration  of 
the  reacting  solutions  that  a  compound  having  the  composition 
given  above  can  be  obtained.  It  is  much  easier  to  be  sure  that 
all  of  the  phosphoric  acid  is  precipitated  from  the  solution  than 
to  secure  the  correct  ratio  between  molybdic  anhydride  and  phos- 
phoric acid.  For  this  reason  it  is  a  common  analytical  practice 
to  dissolve  the  precipitate  in  ammonia  and  precipitate  the  phos- 
phoric acid  as  magnesium  ammonium  phosphate,  MgNH4PO4. 

Tungsten,  W,  184.0.  Although  tungsten  is  usually  classed 
among  the  less  common  elements,  its  compounds  have  been 
known  since  the  middle  of  the  eighteenth  century.  As  a  constit- 
uent of  the  old  "  Damascus  blade  "  its  valuable  effect  on  steel 
was  used  long  before  its  presence  was  recognized,  and  some  of  its 
compounds  have  long  been  used  for  fireproofing  fabrics.  The 
recent  use  for  the  filaments  of  electric  lights  has  now  made  the 
name  of  the  metal  a  household  word. 

Tungsten  is  found  chiefly  in  the  form  of  wolframite,  a  ferrous 
manganese  tungstate,  (FeMn)WO4,  the  iron  and  manganese 
replacing  each  other  as  isomorphous  constituents.  Metallic 
tungsten  can  be  obtained  by  the  reduction  of  tungstic  anhydride, 
WO3,  with  carbon  or  hydrogen  at  a  high  temperature,  or  by 
Goldschmidt's  thermite  process. 

Tungsten  is  a  heavy,  steel-gray,  very  hard  metal.  A  good  deal 
of  difficulty  was  experienced  in  learning  how  to  draw  the  metal 
into  wire  suitable  for  incandescent  electric  lights.  The  specific 
gravity  is  18.64.  Tungsten  melts  at  3000°,  the  highest  melting 
point  of  any  element  except  carbon.  Its  use  for  electric  lights 
depends,  of  course,  on  this  property.  The  temperature  for  rapid 
volatilization  is  probably  higher  than  that  of  carbon. 

A  very  important  application  of  tungsten  is  in  the  manufacture 
of  "  high-speed  tool  steel."  Ordinary  steel  cannot  be  used  for 
any  rapid  lathe  work  because  it  would  become  so  hot  as  to  lose 
its  temper.  Some  kinds  of  tungsten  steel  are  very  hard  and  will 
also  retain  their  hardness  even  when  almost  red-hot.  The 


TUNGSTEN.    URANIUM  531 

introduction   of   such   tools    has   almost    revolutionized   shop 
practice  in  America. 

*  Compounds  of  Tungsten.     The  element  forms  four  chlo- 
rides :  tungsten  dichloride,  WC12,  tungsten  tetrachloride,  WCU, 
tungsten  pentachloride,  WCls,  and  tungsten  hexachloride,  WC16. 
The  last  two  are  volatile,  while  the  dichloride  and  tetrachloride 
are  not.     The  two  most  important  oxides  of  tungsten  are  the  di- 
oxide, WO2,  and  the  trioxide,  or  tungstic  anhydride,  WOa.    Very 
many  complex  tungstates  have  been  prepared.    In  addition  to  the 
normal  sodium  tungstate,  Na2WO4,  no  less  than  thirteen  complex 
sodium  tungstates  containing  a  smaller  proportion  of  tungsten 
have  been  described.     The  compound  5  Na2O.12  WO  3  may  be 
mentioned  as   an  illustration   of   these  compounds.      Sodium 
tungstate  has  been  used  as  a  mordant  in  dyeing  and  for  the 
fireproofing  of  fabrics. 

By  the  reduction  of  acid  sodium  tungstates  by  heating  them 
with  tin  or  hydrogen  a  series  of  yellow,  blue,  violet  and  purple 
compounds  called  "  bronzes  "  has  been  prepared.  These  con- 
tain less  oxygen  than  true  tungstates  should.  Thus  the  composi- 
tion Na2W4Oi2  is  given  for  the  violet  bronze,  while  the  corre- 
sponding tungstate  would  be  Na2W4Oi3  =  Na2O.4  WOs. 

*  Phosphotungstic  Acid,  H3PO4.12  WO3.18  H2O.     This  com- 
pound can  be  prepared  by  treating  silver  tungstate,  Ag2WO4, 
mixed  with  the  calculated  amount  of  phosphoric  acid,  with 
hydrochloric  acid.     It  crystallizes  in  rhombic  crystals  which  are 
soluble  in  water.     The  solution  gives  characteristic  precipitates 
with  alkaloids  and  with  proteins  and  is   used  as  a  reagent  for 
these  purposes.     Several  other  complex  phospho tungstic  acids 
have  been  prepared. 

Uranium,  U,  238.5.  Uranium  is  found  chiefly  in  the  form  of 
uraninite,  or  pitchblende,  or  UaOg.  Several  other  minerals  con- 
tain uranium  and  all  of  these  are  now  of  interest  because  of  the 
connection  with  radium  (p.  471).  Cleveite,  a  mineral  from 
Norway,  resembles  pitchblende,  but  contains  also  yttrium  and 
other  rare  elements.  It  is  the  mineral  in  which  Ramsay  first 
discovered  helium  on  the  earth  (p.  237).  Carnotite  is  a  uranate 


532  A  TEXTBOOK   OF  CHEMISTRY 

and  vanadate  of  potassium  K2O.2  UO3.V2O5.3  H2O.  Samars- 
kite  is  a  complex  tantalo-columbate  of  uranium,  yttrium,  iron 
and  other  metals. 

Uranium  is  a  white  metal.  When  free  from  carbon  it  is  not  so 
hard  as  steel.  Its  specific  gravity  is  18.68.  It  melts  at  a  higher 
temperature  than  platinum. 

The  oxides  of  uranium  are  uranium  dioxide,  UO2,  a  green 
oxide,  U3O8,  of  the  same  composition  as  uraninite  and  uranium 
trioxide  or  uranic  anhydride,  UO3.  The  chlorides  are  uranium 
trichloride,  UC13,  uranium  tetrachloride,  UCU,  uranium  penta- 
chloride,  UCls,  and  uranium  hexachloride,  UCle. 

Comparatively  few  salts  of  uranium  are  known  in  which  metal- 
lic uranium  replaces  the  hydrogen  of  an  acid  directly.  In  these 
few  it  is  quadrivalent.  One  of  the  simplest  is  uranium  sulfate, 
U(SO4) 2- 2  H2O.  The  more  common  salts  of  uranium  contain 
the  bivalent  group  uranyl,  UO2.  Thus  uranyl  nitrate  is 
UO2(NO3)2.3  H2Oand  uranyl  acetate  is  UO2(C2H3O2)2.2  H2O. 

The  compounds  in  which  uranium  acts  as  an  acid  forming  ele- 
ment are  mostly  diuranates,  corresponding  to  the  dichromates. 
Potassium  diuranate,  K2U2O7,  is  an  orange-yellow,  almost  in- 
soluble powder. 

As  a  radioactive  element  uranium  has  a  "  half-life  "  of  about 
6,000,000,000  years. 


CHAPTER  XXXI 
MANGANESE 

Group  VII.  While  there  are  three  elements  (V,  Cb,  Ta)  al- 
ternating with  phosphorus,  arsenic,  antimony  and  bismuth  in 
Group  V,  and  four  elements  (Cr,  Mo,  W,  U)  alternating  with 
sulfur,  selenium  and  tellurium  in  Group  VI,  manganese  is  the 
only  element  alternating  with  chlorine,  bromine  and  iodine  in 
Group  VII.  Not  only  is  the  halogen  of  atomic  weight  about 
214  missing,  but  elements  resembling  manganese  with  atomic 
weights  approximately  98,  187  and  242  have  never  been  found. 
The  radioactivity  of  uranium  and  the  ephemeral  life  of  niton 
(p.  474)  which  are  found  in  this  region  of  the  table  have  recently 
given  us  a  hint  as  to  a  possible  reason  for  these  gaps  in  the  sys- 
tem. It  seems  likely  that  the  structure  of  the  atom  which  would 
give  elements  of  these  atomic  weights  is  unstable,  and  that  either 
these  elements  cannot  exist  at  all  or  they  are  to  be  looked  for 
among  the  radioactive  elements  of  brief  life. 

Manganese  has  properties  such  as  its  position  in  the  table 
leads  us  to  expect.  As  a  metal  it  closely  resembles  iron,  cobalt 
and  nickel,  all  four  elements  forming  a  transition  from  the  hard, 
difficultly  fusible  chromium  to  copper  with  its  much  lower  melt- 
ing point  (1083°)  and  its  great  malleability  and  ductility.  In  its 
nonmetallic  properties,  on  the  other  hand,  it  forms  acidic  oxides 
and  acids,  and  the  highest  of  these,  permanganic  acid,  HMnC>4, 
corresponds  to  perchloric  acid. 

Manganese,  Mn,  54.93.  Occurrence,  Properties.  Manganese 
is  found  chiefly  as  the  dioxide  in  the  mineral  pyrolusite,  MnC>2. 
It  is  also  found  in  small  amounts  in  most  minerals  and  rocks,  in 
practically  all  iron  ores  and  in  some  natural  waters.  Except 
for  scientific  purposes  the  element  is  not  prepared  in  the  pure 
state.  Pure  manganese  is  a  very  hard,  reddish  gray  metal  with 

533 


534  A  TEXTBOOK  OF  CHEMISTRY 

a  specific  gravity  of  7.2.  It  melts  at  1260°.  It  dissolves  easily 
in  acids,  even  more  easily  than  iron,  forming  manganous  salts, 
in  which  the  element  is  bivalent. 

Alloys  of  manganese  with  iron  are  easily  manufactured,  com- 
mercially, by  reducing  a  mixture  of  the  ores  of  the  two  metals 
in  a  blast  furnace  (p.  541).  Those  containing  10-15  per  cent 
of  manganese  are  white,  with  a  brilliant  metallic  luster,  and  re- 
tain the  carbon  in  the  combined  form.  Because  of  the  appear- 
ance of  the  surface,  this  alloy  is  called  spiegeleisen  (mirror-iron). 
It  is  used  in  the  manufacture  of  steel  (p.  548).  An  alloy  con- 
taining 70-90  per  cent  of  manganese  with  iron  is  called  ferro- 
manganese  and  is  used  as  an  addition  to  cast  iron.  Manganese 
bronze  is  an  alloy  of  manganese  and  copper  containing  30  per 
cent  of  manganese.  It  is  hard  and  has  a  high  tensile  strength. 

Compounds  of  Manganese.  In  compounds  in  which  it  acts  as 
a  metal,  manganese  is  almost  exclusively  bivalent,  as  in  manga- 
nous chloride,  MnCl2,  and  manganous  sulfate,  MnSO4.  There 
are  a  few  unstable  compounds,  such  as  manganic  chloride,  MnCU, 
in  which  it  is  trivalent ;  and  compounds  in  which  it  is  a  quadri- 
valent, basic  element  have  been  prepared  only  in  the  form  of 
double  salts  such  as  2  KCLMnCl4,  or  K2MnCl6.  The  basic 
oxides  are  manganous  oxide,  MnO,  and  manganese  sesquioxide, 
Mn2O3. 

As  an  acid-forming  element  manganese  is  quadrivalent  in 
manganese  dioxide,  MnO2,  and  the  manganites,  such  as  calcium 

/°\ 

manganite,  Ca^       y>Mn=O.     It  seems  to  be  sexivalent  in  the 

XX 

K— (X  .0 

manganates,   such  as  potassium  manganate,  /Mnz'   , 

K—<y      ^o 

or  K2MnO4,  and  septivalent  in  the  permanganates,  such  as 

potassium  permanganate,  K — O — Mn=O. 

\) 

Manganous  Manganic  Oxide,  Mn3O4,  is  of  a  mixed  type,  some- 
what similar  to  that  of  red  lead,  but  the  similarity  in  formulas 


MANGANESE  535 


is  probably  superficial.     Red  lead  is  lead  plumbate, 

or  Pk,  while  the  oxide,  Mn3O4,  is  most  likely  derived 


from  a  hypothetical  acid  HMnO2.     If  this  is  true,  the  structure 

xO— Mn=O 
is  Mn(MnO2)2  or  Mn<^ 

\)—Mn=O 

*  Manganous  Hydroxide,  Mn(OH)2,  is  a  white  precipitate 
which  quickly  turns  brown  from  oxidation  on  exposure  to  the  air, 
especially  if  an  alkali  is  present. 

*  Manganous  Chloride,  MnCl2.4  H2O,  is  a  light  pink,  easily 
soluble  salt. 

*  Manganous  Sulfate,  MnSO4,  forms  hydrates  with  1,  2,  3,  4, 
5  or  7  molecules  of  water.     The  last,  which  corresponds  in  com- 
position to  white  vitriol,  ZnSO4.7  H2O,  and  green  vitriol  or 
copperas,  FeSO4.7  H2O,  can  only  be  obtained  by  crystalliza- 
tion at  temperatures  below  6°. 

*  Manganous  Sulfide,  MnS.     Ammonium  sulfide  precipitates 
from  alkaline  solutions  of  manganous  salts  a  flesh-colored  precip- 


itate  having  the  composition  MnS.H2O  or  Mn          .    It  is  easily 

soluble  in  acids,  even  in  acetic  acid.  The  sulfide,  MnS,  is  an 
olive  green  powder  which  is  formed  by  the  action  of  hydrogen 
sulfide  on  any  of  the  oxides. 

Manganese  Dioxide,  MnO2,  or  Black  Oxide  of  Manganese  is 
found  as  the  mineral  pyrolusite  in  sufficient  quantities  to  make  it 
a  very  valuable  commercial  product.  The  compound  has  played 
a  commanding  role  in  the  development  of  chemical  knowledge 
and  also  in  a  variety  of  industrial  processes.  By  means  of  it 
Scheele  discovered  chlorine,  and  it  was  early  used  in  the  prepara- 
tion of  oxygen.  It  is  still  constantly  used  in  laboratories  to 
catalyze  the  decomposition  of  potassium  chlorate  (p.  21). 

During  the  latter  half  of  the  nineteenth  century  manganese 


536  A  TEXTBOOK  OF  CHEMISTRY 

dioxide  was  used  in  large  quantities  for  the  preparation  of  chlo- 
rine as  an  adjunct  to  the  Leblanc  soda  process  (p.  411).  When 
the  supply  of  manganese  ore  declined  and  the  material  grew  ex- 
pensive, the  Weldon  process  (p.  103)  was  introduced  to  conserve 
the  manganese.  At  the  present  time,  when  electrolytic  sodium 
hydroxide  and  chlorine  are  promising  to  complete  the  extinction 
of  Leblanc  soda  and  compete  strongly  with  ammonia-soda,  man- 
ganese is  finding  increasing  use  in  the  metallurgy  of  iron  and 
steel,  and  iron  manufacturers  are  searching  eagerly  for  new  sup- 
plies of  the  ore. 

It  is  worthy  of  notice  that  the  Weldon  process  (p.  103)  depends 
on  the  tendency  of  the  metallic  elements  to  assume  a  higher 
valence  toward  oxygen,  especially  in  an  alkaline  solution,  than 
their  usual  valence  toward  chlorine  and  other  acid  radicals  in 
an  acid  solution.  The  sign  of  the  valence  does  not,  however, 
change.  Manganese  seems  to  be  positive  in  all  of  its  com- 
pounds. 

Manganates.  When  manganese  dioxide,  or  indeed  almost 
any  compound  of  manganese,  is  fused  with  potassium  or  sodium 
carbonate  and  some  oxidizing  agent,  as  potassium  nitrate  or 
potassium  chlorate,  potassium  manganate,  K^MnO^  or  sodium 
manganate,  Na2MnO4,  is  formed.  This  has  a  green  color  and 
dissolves  in  water  to  a  dark  green  solution.  The  solution  is 
sometimes  called  "  chameleon  solution  "  because  of  the  color 
changes  which  it  undergoes  so  easily  as  the  manganate  changes 
to  red  permanganate  or  turns  brown  from  the  separation  of  man- 
ganese dioxide.  The  manganates  are  stable  only  in  an  alkaline 
solution,  and  free  manganic  acid,  H2MnO4,  does  not,  apparently, 
exist,  even  in  solution.  The  addition  of  an  acid,  even  of  carbonic 
acid,  to  the  solution  causes  an  autoxidation  and  reduction  some- 
what similar  to  that  of  hypochlorous  acid  (p.  127)  or  to  that  of 
potassium  chlorate  to  potassium  perchlorate  and  potassium  chlo- 
ride (p.  128). 

H2MnO4  -f  2  H2MnO4  =  2  HMnO4  +  MnO2  +  2  H2O 

Manganic  Permanganic 

Acid  Acid 


PERMANGANATES  537 

(In  developing  this  equation,  notice,  (1)  that  MnO2  +  H2O  is 
equivalent  to  H2MnO3,  hence  one  molecule  of  manganic  acid 
gives  one  atom  of  available  oxygen,  and  (2)  that  one  atom  of 
oxygen  will  oxidize  two  molecules  of  manganic  acid  to  perman- 
ganic acid  by  the  removal  of  two  atoms  of  hydrogen.) 

Permanganates.  Manganic  acid  is  a  very  weak  acid  and  its 
salts  are  strongly  hydrolyzed  in  the  absence  of  an  alkali : 

K2MnO4  +  2  HOH  =  H2MnO4  +  2  KOH 

As  has  just  been  stated,  the  manganic  acid  formed  is  unstable 
and  decomposes  to  manganese  dioxide  and  permanganic  acid, 
HMnO4.  Permanganic  acid  is  a  comparatively  strong  acid  and 
will  at  once  react  with  the  alkali  present,  forming  a  perman- 
ganate. 

Potassium  Permanganate,  KMnO4,  is  a  dark  red,  moderately 
soluble  salt,  which  dissolves  in  water  to  an  intensely  colored  red 
solution.  The  color  is  so  deep  that  the  presence  of  the  perman- 
ganate ion  is  evident  even  in  exceedingly  dilute  solutions.  By 
reduction  in  acid  solutions  the  manganese  is  easily  reduced  to  a 
manganous  salt,  such  as  MnSO4,  which  gives  a  practically  color- 
less solution.  In  alkaline  solutions  the  manganese  is  reduced  to 
manganese  dioxide,  which  is  brown,  leaving  the  solution  above 
colorless.  As  the  loss  of  color  gives  a  very  sharp  indication  of  the 
end  of  the  reaction,  these  properties  are  extensively  used  in 
quantitative  analysis. 

As  typical  reactions  of  this  character  may  be  mentioned  the 
oxidation  of  ferrous  sulfate,  FeSO4,  to  ferric  sulfate,  Fe2(SO4)3, 
in  the  presence  of  sulf uric  acid ;  the  oxidation  of  sulf urous  acid, 
H2SO3,  to  sulfuric  acid ;  of  nitrous  acid,  HNO2,  to  nitric  acid, 
HNOs ;  of  oxalic  acid,  H2C2O4,  to  carbon  dioxide,  CO2 ;  and  of 
hydrogen  peroxide,  H2O2,  to  water,  H2O,  and  oxygen,  O2. 

As  reactions  in  alkaline  solutions  may  be  given  the  oxidation 
of  sodium  sulfite,  Na2SOs,  to  sodium  sulfate,  Na2SO4,  and  that 
of  manganous  sulfate,  MnSO4,  to  manganese  dioxide,  MnO2. 
The  student  is  advised  to  write  the  equations  for  all  of  these 
reactions. 


538  A  TEXTBOOK  OF  CHEMISTRY 

Manganese  Heptoxide,  or  Permanganic  Anhydride,  Mn207, 
is  a  dark  greenish  black,  oily,  volatile  liquid  formed  when  potas- 
sium permanganate  is  treated  with  cold,  concentrated  sulfuric 
acid.  It  is  extremely  unstable,  exploding  violently  on  slight 
provocation.  Many  organic  substances  take  fire  when  brought 
into  contact  with  it. 

It  will  be  seen  on  looking  back  through  the  chapter  that  there 
are  five  well  defined  oxides  of  manganese,  MnO,  Mn3O4,  Mn2O3, 
MnO2  and  Mn2O7.  There  is  some  evidence,  also,  but  not 
entirely  satisfactory,  of  the  existence  of  a  trioxide,  MnO3. 


CHAPTER  XXXII 
IRON,  COBALT,   NICKEL 

Group  VIII.  Between  manganese  and  copper,  molybdenum 
and  silver,  and  between  tungsten  and  gold  there  is,  in  each  case, 
a  group  of  three  metals,  closely  resembling  each  other  in  proper- 
ties. These  three  groups  are :  the  iron  group  —  iron,  cobalt 
and  nickel ;  the  ruthenium  group — ruthenium,  rhodium  and  pal- 
ladium ;  and  the  platinum  group  —  osmium,  iridium  and  plati- 
num. The  valence  of  these  elements  is  variable,  as  with  the 
other  elements  of  this  part  of  the  table.  Valences  of  two,  three 
and  four  are  most  common,  but  nickel  carbonyl,  Ni(CO)4,  ruthe- 
nium oxide,  RuO4,  and  osmium  oxide,  OsO4,  indicate  a  maximum 
valence  of  eight  in  some  cases. 

Iron,  Fe,  55.84.  Iron  is  by  far  the  most  important  of  the 
metals  —  more  important  than  all  of  the  other  metals  taken 
together.  Several  factors  contribute  to  this  importance.  With 
the  exception  of  aluminium,  iron  is  more  abundant  than  any 
other  metal.  As  has  been  pointed  out,  oxygen  forms  about  one 
half  of  the  crust  of  the  earth  and  silicon  one  fourth.  Aluminium 
forms  about  one  fourteenth  and  iron  one  twentieth,  the  four 
elements  comprising  about  seven  eighths  of  that  part  of  the  earth 
which  we  can  examine.  In  addition  to  this,  ores  of  iron  con- 
taining 50  to  70  per  cent  of  the  metal  are  abundant  and  can  be 
reduced  on  a  large  scale  at  a  very  low  cost.  '  Although  alumin- 
ium is  more  abundant  than  iron,  the  latter  can  be  produced 
at  a  very  much  lower  cost.  Finally,  the  forms  of  commercial 
iron  contain  carbon  and  other  elements  which  greatly  change 
its  character,  making  it  possible  to  prepare  many  forms  of  iron 
differing  in  hardness,  malleability,  tensile  strength,  permeability 
to  magnetism  and  other  properties  which  adapt  the  various  forms 
to  special  uses. 

539 


540  A  TEXTBOOK  OF  CHEMISTRY 

Occurrence  of  Iron.  The  ores  of  iron  which  are  of  primary 
value  for  manufacturing  purposes  are  all  of  them  oxides  or  com- 
pounds which  can  be  readily  converted  into  oxides  by  heat. 
The  most  common  and  important  are  hematite  or  ferric  oxide, 
Fe2O3,  magnetite,  a  ferrous-ferric  oxide,  FeaO^  limonite,  ferric 
hydroxide,  Fe2O3.Fe2(OH)6  or  2  Fe2O3.3  H2O,  and  siderite 
or  ferrous  carbonate,  FeCOa.  Iron  pyrites,  FeS2,  is  of  value 
primarily  for  the  sulfur  which  it  contains,  sulfur  being  a  much 
more  expensive  element  than  iron.  The  oxide  of  iron  from  the 
pyrite  burners  of  sulfuric  acid  plants  is  occasionally  used  for  the 
manufacture  of  a  low  grade  of  pig  iron.  The  presence  of  small 
amounts  of  manganese,  phosphorus  and  sulfur  affect  the  quality 
of  the  iron  and  are  often  of  great  importance  in  determining  the 
value  of  an  iron  ore.  Meteorites  often  consist  largely  of  metallic 
iron  usually  containing  nickel,  and  dredgings  from  the  bottom  of 
the  ocean  show  the  presence  of  meteoric  dust  containing  iron. 
The  composition  of  meteorites,  the  presence  of  iron  in  the  sun 
and  density  of  the  earth  as  a  whole  (about  5.53  as  compared  with 
2.7  for  the  portion  we  can  examine,  omitting  the  ocean)  all 
suggest  the  possibility  that  the  central  portions  of  the  earth 
may  contain  large  amounts  of  metallic  iron. 

Metallurgy  of  Iron.  In  prehistoric  times  men  learned  how  to 
reduce  iron  from  its  ores  in  a  forge  or  small  open  hearth,  with  the 
use  of  a  bellows  or  other  device  to  secure  a  blast.  The  process 
was  very  wasteful  of  fuel  and  ore,  but  gave  an  iron  of  fair  quality, 
sometimes  approaching  the  properties  of  steel.  An  apparatus 
somewhat  resembling  a  blast  furnace  seems  to  have  been  invented 
about  the  close  of  the  fifteenth  centirry,  but  for  two  and  a  half 
centuries  the  ore  was  reduced  by  means  of  charcoal.  At  one 
time  this  use  of  charcoal  threatened  to  cause  the  destruction  of 
the  forests  of  England.  In  1735  the  reduction  by  means  of  coal 
was  discovered,  and  the  use  of  coal  and  coke  gradually  displaced 
the  use  of  charcoal,  though  some  charcoal  iron  is  still  made  for 
special  uses,  because  of  its  purity  and  freedom  from  sulfur. 

The  more  important  features  of  a  modern  blast  furnace  for 
the  manufacture  of  iron  are  shown  in  Fig.  106.  A  mixture  of 


IRON:  BLAST  FURNACE 


541 


ore,  fuel  (usually  coke)  and  limestone  is  introduced  at  the  top 
of  the  furnace  in  such  proportions  that  the  oxides  of  iron  are 
completely  reduced,  giving  metallic  iron.  The  iron  combines 
with  carbon,  silicon  and 
small  amounts  of  phosphorus 
and  sulfur  to  form  the  crude 
product  known  as  pig  iron. 
The  lime  of  the  limestone 
combines  with  the  silica  and 
other  impurities  of  the  ore, 
forming  a  fusible  silicate, 
called  a  slag,  which  melts 
and  collects  in  the  bottom 
of  the  furnace  on  top  of  the 
melted  iron.  The  air  for  the 
combustion  of  the  fuel  is 
forced  in  by  means  of  a 
powerful  blower  through  the 
openings  near  the  bottom  of 
the  furnace,  which  are  called 
tuyeres.  The  fuel  remains 
in  excess  down  to  the  very 
bottom  of  the  furnace,  with 
the  result  that  carbon  diox- 
ide and  water,  formed  as  the 
ore  is  reduced,  are  continu- 
ally reduced  back  to  carbon 
monoxide  and  hydrogen. 
The  reduction  is  brought 
about  chiefly  by  the  carbon 
monoxide  and  hydrogen  in 


accordance     with     the 
versible  reactions : 


re- 


Fig.  106 


Fe2O3  +  3  CO  :±  2  Fe  +  3  CO2 


and 


Fe2O3  +  3  H2:z±2  Fe  +  3  H2O 


542  A  TEXTBOOK  OF  CHEMISTRY 

The  equilibrium  of  these  reactions  is  very  far  toward  the  left, 
and  it  is  only  because  the  solid  carbon  of  the  fuel  continually 
reduces  the  carbon  dioxide  and  water  back  to  carbon  monoxide 
and  hydrogen  that  the  process  can  succeed.  These  conditions 
make  it  necessary  to  use  such  a  proportion  of  fuel  that  the  gases 
escaping  from  the  top  of  the  furnace  still  contain  considerable 
amounts  of  carbon  monoxide  and  hydrogen  —  enough  so  that 
these  blast  furnace  gases  furnish  a  valuable  fuel,  retaining  ap- 
proximately one  half  of  the  original  energy  of  the  coke.  The 
gases  are  used  to  furnish  the  power  for  the  blowing  engines, 
either  by  burning  them  under  boilers  used  to  furnish  steam  for 
engines  or  by  utilizing  them  directly  in  gas  engines.  The  gases 
are  also  used  in  "  stoves  "  in  which  the  blast  of  air  for  the  furnace 
is  heated  to  about  800°  before  it  enters  the  tuyeres.  The  use  of 
the  hot  blast  concentrates  the  reactions  of  the  furnace  in  the  lower 
part  and  greatly  lessens  the  amount  of  fuel  required  in  the  charge. 
It  has  been  discovered  rather  recently  that  a  further  saving  of 
about  10  per  cent  in  the  amount  of  coke  required  per  ton  of 
iron  is  effected  by  first  cooling  the  air  for  the  blast  to  a  low  tem- 
perature so  as  to  condense  most  of  the  moisture  which  it  contains 
(Gayley ;  see  Journal  of  Industrial  and  Engineering  Chemistry, 
5,  241  (1913)). 

The  materials  used  in  a  furnace  working  with  an  ore  containing 
60  per  cent  of  iron  are,  approximately,  in  the  proportion,  one  ton 
of  ore,  0.6  ton  of  coke  and  0.3  ton  of  limestone,  but  the  amounts 
vary  with  the  character  of  the  gangue  in  the  ore  and  the  ash  of 
the  coke.  Approximately  five  tons  of  air  must  be  blown  into 
the  furnace  for  each  ton  of  ore  reduced.  The  process  is  carried 
on  continuously  for  many  months  and  sometimes  for  several 
years.  At  intervals  of  a  few  hours  the  melted  iron  which  col- 
lects in  the  hearth  of  the  furnace  is  drawn  off  through  the  tap- 
hole  at  the  bottom  into  a  large  ladle,  by  means  of  which  it  is 
transferred  to  a  mixer,  where  the  iron  from  several  furnaces  is 
brought  together  and  mixed  before  treating  it  further  in  Besse- 
mer converters  or  open  hearth  furnaces.  The  larger  part  of  the 
iron  is  not  allowed  to  cool  until  it  is  converted  into  steel  rails, 


CAST   IRON  543 

steel  or  iron  plates  or  structural  materials  of  various  forms. 
In  the  older  practice,  in  the  production  of  pig  iron  for  foundry 
use  or  for  the  market,  the  iron  was  drawn  out  into  a  series  of 
channels  in  the  sand  floor  of  the  furnace  room  and  allowed  to 
solidify.  It  is  then  called  pig  iron.  In  modern  practice  pig 
iron  is  usually  cast  for  the  market  in  continuous  casting  machines. 

The  slag  either  runs  continuously  from  an  opening  above  the 
surface  of  the  iron  in  the  hearth  or  more  often  is  drawn  off  each 
time  after  the  iron.  The  furnace  slag  is  now  extensively  used 
for  the  manufacture  of  Portland  cement  and  has  sometimes  been 
used  to  make  cheap  glass. 

Pig  Iron.  Cast  Iron.  The  iron  from  the  blast  furnace  is 
always  a  crude  product  containing  manganese,  carbon,  partly 
combined  with  the  iron,  partly  as  graphite,  silicon,  sulfur  and 
phosphorus.  The  larger  portion  of  this  crude  iron  is  subjected 
to  various  methods  of  treatment  which  convert  it  into  steel  or 
refined  irons  which  are  more  suitable  for  most  purposes  than 
the  crude  iron.  A  large  amount  of  pig  iron,  however,  is  melted 
in  cupola  furnaces  with  or  without  the  addition  of  aluminium, 
ferromanganese  or  other  substances  to  improve  its  quality,  and 
cast  in  sand  moulds.  Iron  prepared  is  this  way  is  gray  in  struc- 
ture. In  the  melted  iron  the  carbon  is  probably  all  combined 
with  the  iron  as  iron  carbide,  FesC,  which  dissolves  in  the 
molten  mass,  forming  a  homogeneous  solution.  When  the  iron 
is  cast  as  described,  however,  the  larger  part  of  the  carbon  sepa- 
rates from  the  iron  as  graphite  during  the  slow  cooling.  Such 
an  iron  has  a  gray  color  and  is  known  as  gray  cast  iron.  The 
presence  of  the  graphite  can  be  easily  shown  by  dissolving  the 
iron  in  a  dilute  acid. 

If  the  iron  is  cast  in  contact  with  a  cold  metallic  surface,  called 
a  "  chill,"  the  carbon  does  not  have  time  to  separate  as  graphite, 
but  remains  combined  with  the  iron,  giving  a  very  hard  white 
iron,  suitable  for  the  rims  of  car  wheels  and  other  similar  pur- 
poses. The  addition  of  ferromanganese  aids  in  holding  the 
carbon  in  the  combined  form.  On  the  other  hand,  silicon  tends 
to  cause  the  carbon  to  separate  as  graphite. 


544 


A  TEXTBOOK  OF  CHEMISTRY 


The  following  analyses  indicate  the  usual  composition  of  the 
material : 1 

ANALYSES  OF  PIG  IRON 


l 

2 

Iron  (by  difference)       

9429 

92  72 

0.55 

061 

Graphite      

2.22 

1  85 

Silicon     

1  84 

2  57 

0035 

0044 

Phosphorus      

0.19 

0.54 

Titanium     

0.074 

0081 

Manganese      . 

074 

1  54 

Copper   

0.06 

0.043 

100.00 

100.00 

Gray  cast  iron  melts  at  1120°  to  1230°. 

Wrought  Iron.  While  castings  of  gray  iron  can  be  finished 
by  filing  or  turning  in  a  lathe,  the  metal  cannot  be  welded  or 
rolled  into  bars  or  sheets.  The  puddling  process  for  producing 
a  nearly  pure  iron  from  the  cast  iron  was  invented  by  Henry 
Cort  in  England  in  1784.  The  iron  is  melted  and  subjected  to 
an  oxidizing  flame  on  a  hearth  lined  with  iron  ore.  It  oxidizes 
to  the  magnetic  oxide,  Fe3O4,  on  the  surface,  and  by  stirring 
this  through  the  mass  of  the  iron  with  a  rabble  the  oxygen  of 
the  oxide  burns  the  carbon  to  carbon  monoxide,  CO,  which 
escapes.  The  silicon  is  burned  to  silicon  dioxide,  which  com- 
bines with  ferrous  oxide,  FeO,  forming  a  fusible  ferrous  silicate, 
Fe2SiO4.  The  phosphorus  is  burned  to  phosphorus  pentoxide, 
P2O5,  which  combines  with  more  ferrous  oxide  to  ferrous  phos- 
phate, Fes(PO4)2,  and  the  sulfur  burns  to  sulfur  dioxide,  which 
escapes.  The  process  gives  a  malleable,  ductile  iron  which 

1  Samples  of  pig  iron  furnished  by  the  U.  S.  Bureau  of  Standards. 
The  analyses  are  published  with  the  permission  of  the  Director  of 
the  Bureau. 


STEEL  545 

may  be  more  than  99  per  cent  pure  iron.  This  process,  which 
was  a  very  important  one  till  near  the  close  of  the  nineteenth 
century,  has  now  been  largely  replaced  by  the  Bessemer  and 
Open  Hearth  processes,  which  are  used  to  make  mild  irons  as 
well  as  steel.  Pure  iron  melts  at  1530°  and  has  a  specific  gravity 
of  7.55.  It  boils  at  2950°. 

Cementation  Steel.  Cast  Steel.  Wrought  iron  contains  only 
a  very  small  per  cent  of  carbon.  It  is  soft  and  malleable  and 
can  be  welded,,  but  will  not  harden  and  is  not  suitable  for  the 
manufacture  of  knives  and  edge  tools.  By  packing  bars  of 
wrought  iron  in  charcoal,  in  long,  earthenware  boxes  and  heating 
them  to  1000°-!  100°  for  8  to  10  days  the  iron  absorbs  approxi- 
mately one  per  cent  of  carbon  and  is  changed  to  steel.  The 
process  is  called  cementation.  The  product  is  melted  to  render 
it  homogeneous  and  is  then  known  as  cast  steel.  This  process, 
which  was  practically  the  only  method  known  for  making  a 
good  grade  of  steel  before  1855,  has  been  more  and  more  dis- 
placed by  the  Bessemer  and  open  hearth  processes.  The  cemen- 
tation process  is  tedious  and  very  expensive,  and  is  now  used 
only  for  the  manufacture  of  a  very  high  grade  of  steel  for  mak- 
ing cutlery  and  for  other  uses  where  the  articles  manufactured 
are  small  and  the  labor  expended  in  giving  them  their  final  form 
is  the  chief  item  in  the  cost  of  production.  The  name  "  cast 
steel "  is  now  often  used  for  steels  made  by  other  methods. 

When  steel  is  heated  to  a  temperature  of  bright  redness, 
700°-800°,  and  then  cooled  suddenly  by  quenching  in  water, 
it  is  rendered  very  hard.  If  such  a  steel  is  heated  again  to  a 
temperature  of  450°-600°  the  steel  becomes  less  hard  and  less 
brittle.  This  process  is  called  tempering.  By  polishing  the 
steel  before  the  second  heating,  the  color  of  the  film  of  oxide 
which  forms  on  the  surface  furnishes  an  indication  as  to  when 
the  proper  temper  has  been  obtained.  Skilled  workmen  deter- 
mine the  temper  by  watching  the  color  of  the  surface  of  the 
steel. 

*  The  tempering  of  steel  can  be  best  understood  in  the  light 
of  the  following  facts : 


546 


A  TEXTBOOK  OF  CHEMISTRY 


1.  Pure  iron  exists  in  three  allotropic  modifications  called 
a-,  /?-  and  y-  ferrite.  The  word  ferrite  is  simply  a  name  used  by 
metallurgists  to  designate  pure  iron.  These  forms  of  iron 
differ  in  properties,  especially  in  being  magnetic  or  nonmagnetic 
and  in  the  amount  of  iron  carbide,  FesC,  which  they  can  hold 
in  solid  solution.  Each  is  stable  through  a  definite  range  of 
temperature.  The  ranges  of  stability  and  properties  will  be 
seen  from  the  following  table : 


ALLOTROPIC 
FORM 

TEMPERATURE 
OP  STABILITY 

MAGNETIC 
PROPERTIES 

HARDNESS 

DUCTILITY 

SOLUBILITY 
OF  FesC 

a-Ferrite   .     . 
/3-Ferrite  .    . 
•y-Ferrite  .     . 

Below  750° 
750°-860° 
Above  860° 

Magnetic 
Nonmagnetic 
Nonmagnetic 

Soft 
Hard 
Hard 

Ductile 
Brittle 
Ductile 

Little 
Little 
Considerable 

The  existence  of  these  three  forms  has  been  demonstrated  by 
a  study  of  the  rate  of  cooling  of  iron.  If  a  piece  of  iron  which 
has  been  heated  to  a  temperature  above  860°  is  connected  with 
a  thermocouple  to  record  the  temperature,  it  will  be  found  that, 
instead  of  a  regular  fall  in  temperature,  as  would  be  expected, 
when  a  transition  point  is  reached  the  temperature  falls  more 
slowly,  ceases  to  fall,  or  may  even  rise  for  a  short  time,  because 
of  the  heat  evolved  when  y-ferrite  changes  to  /3-ferrite  or  when 
^-ferrite  changes  to  a-ferrite.  In  the  case  of  steel  a  visible 
brightening  of  the  mass,  which  is  called  recalescence,  can  be 
observed  when  one  of  the  transition  points  is  reached. 

2.  Iron  combines  with  carbon  to  form  a  definite  compound, 
iron  carbide,  FeaC,  which  is  called  by  the  metallurgists  cement- 
ite.     This  carbide  is  soluble  in  y-ferrite  but  much  less  soluble 
in  a-ferrite  or  /?-ferrite. 

3.  The  presence  of  the  iron  carbide,  FesC,  lowers  the  tran- 
sition points  so  that  in  a  steel  containing  0.89  per  cent  of  carbon 
this  may  all  be  held  in  a  homogeneous  solution  in  the  y-ferrite 
at  a  temperature  of  690°,  or  above.     Such  a  saturated  solution 
of  cementite  or  iron  carbide  in  y-ferrite  is  called  austenite.     It 
is  very  hard,  and  a  hardened  steel  of  this  composition  consists 


BESSEMER  STEEL 


547 


entirely  of  austenite.  The  quenching  of  the  steel  in  water 
carries  it  so  quickly  by  the  transition  points  that  the  transforma- 
tion does  not  occur  and  the  metal  is  left  in  the  hardened  form. 
4.  If  such  a  steel  is  cooled  slowly  below  the  transition  point 
to  a-ferrite,  the  latter  can  no  longer  hold  the  cementite  or 
carbide  in  solution,  and  the  two  separate  into  a  mixture,  which 
would  consist,  when  the  separation  is  complete,  of  about  13 
per  cent  of  cementite  and  87  per  cent  of  a-ferrite.  As  the  latter 
is  soft  and  forms  the  larger  part  of  the  mass,  it  gives  character 
to  the  whole.  On  the  other  hand,  if  the  hardened  steel  is  heated 
to  450°-600°,  it  changes  slowly  to  the  mixture  of  a-ferrite  and 
cementite,  the  change  being  more  rapid  and  complete  at  the 
higher  temperature.  The  more  completely  the  change  occurs, 
the  softer  will  be  the  steel. 

Bessemer  Steel.  In  1852  an  American  by  the  name  of 
Kelly  patented  a  process  for  purifying  iron  by  blowing  air 
through  it.  Three 
years  later,  Besse- 
mer in  England 
discovered  inde- 
pendently and  pat- 
ented a  similar 
process  and  suc- 
ceeded in  develop- 
ing it  to  practical 
success.  After 
some  litigation 
Mr.  Kelly  sold  out 
his  interest  to  Bes- 
semer, and  the 
process  is  called  by 
the  latter's  name. 
The  apparatus 
used  is  shown  in 
Fig.  107.  The 
large,  cylindrical 


Fig.  107 


548  A  TEXTBOOK  OF  CHEMISTRY 

vessel  is  at  first  turned  on  its  side  and  a  charge  of  several 
tons  of  pig  iron  introduced.  A  strong  blast,  which  enters 
through  one  axis  of  the  converter  and  through  the  tuyere 
holes  at  the  bottom,  is  turned  on  and  the  converter  brought 
to  an  upright  position.  The  silicon  of  the  iron  is  burned 
to  silicon  dioxide  and  the  carbon  to  carbon  monoxide.  The 
silica  combines  with  ferrous  oxide  to  form  a  highly  silicious 
slag.  The  heat  from  the  combustion  of  the  silicon  raises  the 
temperature  of  the  iron  so  that  it  remains  fluid  even  after  the 
carbon  and  silicon  are  removed,  although  the  melting  point  of 
the  iron  is  considerably  raised.  When  the  carbon  is  gone,  the 
flame  suddenly  drops  and  at  this  point  the  converter  is  again 
turned  on  its  side  and  enough  spiegeleisen  added  to  give  a  steel 
containing  the  desired  amount  of  carbon  —  0.40  to  0.45  per 
cent  for  steel  rails.  After  mixing,  the  steel  is  poured  into  ingot 
moulds  and  from  these  it  is  placed  in  "  soaking  pits  "  to  come  to 
a  uniform  temperature  throughout  and  then  taken  directly  to 
the  rolls  and  rolled  into  rails  or  structural  iron  or  other  forms  of 
iron,  without  being  allowed  to  cool. 

In  the  "  acid  "  Bessemer  process  as  described,  in  which  the 
lining  of  the  converter  is  of  silicious  materials,  the  slag  contains 
more  than  70  per  cent  of  silica,  SiO2,  and  very  little  of  the  phos- 
phorus is  removed.  A  good  quality  of  steel  can  be  obtained 
only  when  the  material  used  is  quite  pure.  To  make  use  of 
less  pure  ores  and  iron,  the  basic  or  Thomas-Gilchrist  process 
was  designed.  For  this  the  converter  is  lined  with  calcined 
dolomite  and  lime  is  added  to  the  charge.  In  the  presence  of 
the  basic  lining  and  lime  the  phosphorus  is  removed  as  calcium 
or  magnesium  phosphate.  The  basic  slag  from  the  converter 
is  valuable  as  a  fertilizer  because  of  the  phosphorus  which  it 
contains.  In  America  this  process  has  been  completely  dis- 
placed by  the  basic  open  hearth  process. 

Open  Hearth  or  Siemens-Martin  Process.  During  the  last 
twenty  years  another  process,  illustrated  in  Figs.  108,  109  and 
110,  has  grown  rapidly  in  favor  and  at  the  present  time  more 
steel  is  manufactured  in  the  United  States  by  the  open  hearth 


OPEN  HEARTH  STEEL 


549 


550  A  TEXTBOOK  OF  CHEMISTRY 

than  by  the  Bessemer  process.  In  Fig.  108  two  chambers  are 
represented  filled  with  a  checkerwork  of  brick.  The  producer 
gas  (p.  297)  and  air  used  in  the  furnace  pass  up  through  one  pair 
of  these  chambers,  and  after  passing  through  the  furnace  room 
pass  downward  through  a  second  pair  of  chambers,  parting  with 
their  heat  to  the  bricks.  After  an  interval  of  twenty  minutes 
or  half  an  hour,  the  current  is  reversed,  and  now  the  gas  and 
air  take  up  heat  from  the  bricks  before  entering  the  furnace 
proper.  Such  a  furnace  is  called  a  regenerative  furnace  and 
works  economically  with  a  low  grade  of  gas.  In  the  furnace 
there  is  melted  a  mixture  of  cast  iron,  ore,  steel  scraps  and 
sometimes  lime  or  other  fluxes.  A  charge  of  50  to  75  tons  may 
be  used,  and  as  the  materials  can  be  kept  melted  indefinitely,  it 
is  possible  to  take  out  a  sample  for  analysis  and  secure  a  more 
accurate  control  than  with  the  Bessemer  converter.  The 
process  can  also  utilize  an  almost  unlimited  variety  of  material, 
and,  especially,  it  is  suitable  for  iron  containing  too  much  phos- 
phorus for  the  acid  Bessemer  and  too  little  for  the  basic  Besse- 
mer process.  It  is  possible  to  stop  the  decarburization  of  the 
iron  at  any  point  desired  instead  of  carrying  it  to  completion 
and  recarburizing,  as  is  done  in  the  Bessemer  process.  By  all 
three  processes  grades  of  iron  and  steel  containing  from  0.1  to 
1.0  per  cent  of  carbon  are  made. 

Nails  and  sheet  iron  made  from  materials  produced  by  these 
modern  processes  corrode,  when  subjected  to  moisture  and  air, 
very  much  more  rapidly  than  when  made  from  wrought  iron 
from  the  puddling  process.  This  seems  to  be  because  the  latter 
is  more  homogeneous  and  does  not  give  differences  of  potential 
between  different  parts  of  the  iron.  It  has  recently  been  found 
that  the  addition  of  a  very  small  amount  of  copper  (0.2  per  cent, 
or  less;  Chamberlain  J.  Ind.  and  Eng.  Chem.  5,  360  (1913)) 
renders  the  iron  very  much  more  resistant  to  corrosion.  The 
effect  is  similar  to  that  produced  by  amalgamating  the  surface 
of  the  zinc  used  in  an  electrical  cell  (p.  481). 

The  following  analyses  illustrate  the  composition  of  iron  and 
steel  made  by  modern  processes. 


IRON  AND  STEEL 


551 


. 


!>•    i—  I 


ss 

CD  CD 


O5   CO  iO 
O   CO  O 


LO  O  O 

LO  O  1— I  T-H 

O  iO  r-H  O 

CO  O  O  O 


CO   ^     ^ 

O    O    _M    CO    OS 


O  t^   T— i   <M 

r^  o  CD  o 

o 


CO   CO  t— 

•^t1  CO  »— i  *O 
CD  iO  Ol  CO 
000^ 


iO  »O  O  <M 
r^  CO  CO  O 
GO  O  O  O 


o  o  o  o 


CS|  GO  O  rH 
O  CO  O  O 
O  O  O  O 


T-i  co 

CO    GO 

o  o 


LO   CO   Oq   CO 
O    CO    r-t    CO 

T-H     CO     O     O 

0000 


c^  o 

o  o 


0000 
O 


00 


O  O  iO   CO 

O  GO  i-H    O 

GO  O  O  O 

o 


SLO 
I  — 
rH    !>. 

o  o 


O    CO    GO   CO 
GO   O   O  O 


<M   CO 

ss 

o  o 


8 


11  „•  J*.M|  • 

111 I 111 I 11 


-a 
$ 


552  A  TEXTBOOK  OF  CHEMISTRY 

Alloy  Steels.  The  addition  of  chromium,  manganese,  tungsten 
and  molybdenum  to  carbon  steels  lowers  the  point  of  decom- 
position of  austenite  (p.  546).  By  the  use  of  suitable  mixtures 
the  point  of  decomposition  may  be  brought  below  ordinary  tem- 
perature, and  such  a  steel  will  be  hard  at  any  temperature. 
Steels  of  this  type  have  proved  very  useful  for  high-speed  lathe 
tools,  which  can  be  run  at  such  a  rate  as  to  become  almost  red- 
hot  without  losing  their  temper.  Such  steels  are  called  "  self- 
hardening." 

Compounds  of  Iron.  When  iron  is  dissolved  in  dilute  acids, 
it  forms  ferrous  salts,  such  as  ferrous  chloride,  FeCU,  or  ferrous 
sulfate,  FeSC>4.  In  these  compounds  the  iron  is  apparently 
bivalent,  but  the  vapor  density  of  ferrous  chloride  at  tempera- 
tures slightly  above  its  boiling  point  points  to  a  formula,  Fe2Cl4, 
rather  than  FeCl2.  If  we  assume  this  as  the  true  formula,  it 
still  remains  uncertain  whether  the  iron  or  the  chlorine  is  tri- 
valent  or  the  iron,  possibly,  quadrivalent.  The  structure  might 
be 

Ck  /Cl  /Cl=Ck 

>Fe=F<  or          Fe<  >Fe 

CK  X3  \C1=CK 

At  present  no  means  has  been  discovered  of  deciding  posi- 
tively between  these  formulas. 

On  exposure  to  the  air,  especially  in  neutral  or  alkaline  solu- 
tions, ferrous  compounds  are  oxidized  to  the  ferric  condition. 
If  an  acid  is  present,  a  ferric  salt  is  formed.  If  the  solution  is 
alkaline,  neutral,  or  acid  with  a  weak  acid,  such  as  carbonic 
acid,  ferric  hydroxide  or  a  basic  salt  is  formed.  Iron  is  ap- 
parently trivalent  in  the  ferric  salts,  but  the  vapor  density  indi- 
cates that  ferric  chloride  has,  in  part,  the  formula  Fe2Cle  at 
temperatures  slightly  above  the  point  of  sublimation.  As  with 
the  ferrous  salts,  it  is  not  known  whether  the  molecule  is  held 
together  by  the  iron  or  the  chlorine  atoms. 

The  basic  properties  of  iron  are  much  weaker  in  the  ferric 
than  in  the  ferrous  salts  and  while  there  is  some  hydrolysis  in 
solutions  of  ferrous  salts  and  the  ferrous  salts  of  strong  acids 


IRON  COMPOUNDS  553 

react  acid  toward  litmus,  the  hydrolysis  of  ferric  salts  is  much 
more  marked.  This  is  well  illustrated  by  the  complete  precipi- 
tation of  the  iron  as  ferric  hydroxide  when  powdered  barium 
carbonate,  suspended  in  water,  is  added  to  a  solution  of  ferric 
chloride : 

FeCl3  +  3  HOH  ^  Fe(OH)3  +  3  HC1 

2  HC1  +  BaCO3  =  BaCl2  +  H2O  +  CO2 

The  ferric  hydroxide  or  a  basic  ferric  salt,  such  as  FeCl2OH 
or  FeCl(OH)2,  remains  in  collodial  solution  until  the  barium 
carbonate  is  added.  When  the  barium  carbonate  reacts  with 
the  hydrochloric  acid,  the  hydrolysis  becomes  complete  and  the 
barium  chloride  also  assists  in  coagulating  the  colloidal  ferric 
hydroxide  (p.  362). 

Potassium  Ferrate,  K2FeO4,  can  be  prepared  by  passing 
chlorine  through  a  solution  of  potassium  hydroxide  in  which 
ferric  hydroxide  is  suspended.  This  gives  a  red  solution  from 
which  the  salt  can  be  crystallized.  Some  other  ferrates  have 
been  prepared,  but  all  of  these  hydrolyze  in  water  and  decompose 
easily. 

Ferrous  chloride,  FeCl2.4  H2O,  separates  from  concentrated 
solutions  from  which  the  air  has  been  carefully  excluded  in 
clear  blue  crystals,  which  become  green  on  exposure  to  the  air. 
The  hydrate  dissolves  in  two  thirds  of  its  weight  of  water.  In 
the  presence  of  hydrochloric  acid  it  is  easily  oxidized  to  ferric 
chloride,  FeCl3,  by  the  action  of  nitric  acid,  potassium  perman- 
ganate, potassium  dichromate,  chlorine  or  almost  any  vigorous 
oxidizing  agent. 

Ferrous  Hydroxide,  Fe(OH)2,  forms  as  a  pure  white  precipi- 
tate on  precipitation  of  a  ferrous  salt  with  sodium  hydroxide  in 
a  solution  entirely  free  from  dissolved  oxygen  or  an  oxidizing 
agent.  The  slightest  exposure  to  the  air  causes  the  precipitate 
to  turn  green,  and  on  longer  exposure  it  is  changed  to  reddish 
brown,  ferric  hydroxide,  Fe(OH)3. 

*  Ferrous  Oxide,  FeO,  is  formed  as  a  black  powder  when 
ferrous  oxalate,  FeC2O4,  is  heated  out  of  contact  with  the  air. 


554  A  TEXTBOOK  OF  CHEMISTRY 

In  some  conditions  it  takes  fire  spontaneously  on  exposure  to 
the  air.  If  reduced  by  hydrogen  at  a  low  temperature  it  forms 
pyrophoric  iron,  which  takes  fire  in  the  air. 

Ferrous  Sulfate,  Green  Vitriol  or  Copperas,  FeSO4.7  H2O. 
Metallic  iron  dissolves  readily  in  dilute  sulfuric  acid  and  the 
hydrate  known  as  green  vitriol  crystallizes  from  the  solution 
on  cooling.  One  hundred  parts  of  water  dissolves  38  parts  of 
the  hydrate  at  10°  or  48  parts  at  20°.  Because  the  salt  is  much 
less  soluble  than  the  chloride,  sulfuric  acid  is  not  so  suitable  as 
hydrochloric  acid  for  the  preparation  of  hydrogen  sulfide  from 
ferrous  sulfide,  FeS. 

Ferrous  Carbonate,  FeCOs,  is  found  in  nature  in  the  mineral 
siderite.  It  crystallizes  in  rhombohedra  and  is  isomorphous 
with  calcite,  CaCOs.  An  impure  siderite,  called  clay  iron  stone, 
has  been  one  of  the  most  important  iron  ores  used  in  England. 

Ferrous  carbonate  dissolves  as  ferrous  bicarbonate,  FeH2  (003)2, 
in  waters  which  contain  carbonic  acid,  exactly  as  calcium  car- 
bonate does.  Such  waters  are  known  as  chalybeate  waters,  and 
some  of  these  have  been  considered  valuable  for  their  medicinal 
properties.  On  exposure  to  the  air  the  iron  of  such  waters  is 
oxidized  to  the  ferric  state  and  separates  as  ferric  hydroxide. 
Beds  of  iron  ore  were  doubtless  formed,  in  many  cases  by  a 
similar  process. 

Ferrous  Chloride  and  Nitric  Oxide.  Solutions  of  ferrous 
chloride  or  ferrous  sulfate  absorb  nitric  oxide  readily,  giving  dark 
brown  or  black  solutions.  As  one  molecule  of  the  ferrous  salt 
will  absorb  one  molecule  of  nitric  oxide,  it  is  supposed  that  the 
solution  contains  a  compound  of  the  formula  FeCl2.NO.nH2O,1 
but  it  has  never  been  possible  to  isolate  the  compound,  and  it  is 
completely  decomposed  by  boiling  the  solution. 

Ferric  Chloride,  FeCl3.6  H2O,  can  be  obtained  by  passing 
chlorine  into  a  solution  of  ferrous  chloride  and  crystallizing  the 
solution.  It  is  a  yellow,  deliquescent  solid,  very  easily  soluble 
in  water.  The  anhydrous  chloride,  FeCls,  sublimes  in  dark 

1  Manchot  and  Zechentmayer,  Liebig's  Annalen,  350,  368  (1906). 


IRON  COMPOUNDS  555 

green,  iridescent  scales  which  are  red  by  transmitted  light.  It 
boils,  or  sublimes,  at  280°-285°.  At  448°  the  weight  of  the  gram 
molecular  volume  is  about  303  grams  and  at  750°  it  is  167  grams. 
As  the  molecular  weight  of  Feds  is  162.2,  it  seems  that  the  mole- 
cules at  the  lower  temperature  are  chiefly  Fe2Cle  and  at  the 
higher  temperature,  FeCls. 

Ferric  chloride  is  easily  reduced  to  ferrous  chloride  by  hydro- 
gen sulfide,  by  nascent  hydrogen,  or  by  stannous  chloride,  SnCl2. 

Ferric  Hydroxide,  Fe(OH)3,  is  readily  formed  as  a  reddish 
brown,  flocculent,  amorphous  precipitate  when  sodium  hydroxide 
or  ammonia  is  added  to  a  solution  of  a  ferric  salt.  It  does  not 
dissolve  in  an  excess  of  the  alkali,  differing  very  markedly  from 
aluminium  hydroxide,  A1(OH)3,  in  this  regard.  Ferric  hydroxide 
loses  water  very  easily,  and  all  of  the  hydroxides  in  nature  con- 
tain less  water  than  would  correspond  to  the  formula  Fe(OH)3. 
The  most  common  natural  hydroxide  is  limonite, 
Fe2O3.Fe2(OH)6,  or  as  it  is  often  written,  2  Fe2O3.3  H2O.  Com- 
pounds of  this  type  illustrate  the  somewhat  artificial  character 
of  the  distinction  between  hydroxides  and  hydrates. 

Dialyzed  Iron.  A  solution  of  ferric  chloride  will  dissolve  a 
considerable  quantity  of  ferric  hydroxide.  If  the  solution  is 
dialyzed  with  a  parchment  membrane  (p.  353),  the  hydrochloric 
acid  formed  by  the  hydrolysis  of  the  chloride  passes  through 
the  membrane,  while  the  ferric  hydroxide  remains  behind  as  a 
colloidal  solution.  Such  a  solution  is  called  "  dialyzed  iron  " 
and  is  sometimes  used  in  medicine.  It  is  particularly  suitable 
as  an  antidote  for  arsenic  poisoning.  The  arsenious  oxide  com- 
bines with  the  ferric  hydroxide  to  form  an  insoluble  arsenite. 

Ferric  Oxide,  Fe2O3,  is  prepared,  artificially,  by  igniting  the 
hydroxide.  It  is  also  found  in  nature  as  the  mineral  hematite. 
Ferric  oxide  prepared  in  various  ways  is  sold  as  a  pigment  under 
the  name  Venetian  red  or  as  rouge.  A  mixture  of  calcium  sulfate 
and  ferric  oxide  prepared  by  calcination  of  a  mixture  of  sulfate 
of  iron  and  lime  is  sometimes  sold  under  the  same  name.  The 
best  quality  of  rouge  is  obtained  by  calcining  ferrous  oxalate. 
The  shade  obtained  depends  on  the  temperature  of  calcination. 


556  A  TEXTBOOK  OF  CHEMISTRY 

*  Ferric  Sulfate,  Fe2  (804)3,  is  formed  by  the  oxidation  of 
ferrous  sulfate  in  the  presence  of  sulfuric  acid.     It  forms  alums, 
of  which  ferric  ammonium  alum,  NH4Fe(SC>4)2.12  H^O,  is  the 
best  known.     Ferric  sulfate  was  formerly  obtained  by  the  oxi- 
dation of  a  pyrite-bearing  shale  at  Nordhausen  in  Bohemia, 
and  was  decomposed  by  dry  distillation  for  the  preparation  of 
fuming  sulfuric  acid. 

Magnetic  Oxide  of  Iron,  FesC^,  is  the  product  formed  when 
iron  burns  in  air  or  oxygen  or  when  steam  is  passed  over  heated 
iron.  It  occurs  also  as  the  natural  mineral  magnetite,  which 
is  one  of  the  purest  of  the  iron  ores  in  Norway  and  Sweden. 
It  may  be  considered  as  ferrous  ferrite,  Fe(FeO2)2,  or  FeO.Fe2Os, 
formulas  which  bring  out  its  relation  to  chromite,  with  which  it 
is  isomorphous. 

Ferrous  Sulfide,  FeS,  is  readily  formed  by  heating  a  mixture 
of  iron  and  sulfur.  It  may  also  be  prepared  by  heating  iron 
pyrites,  FeS2,  which  loses  half  of  its  sulfur  at  a  high  temperature. 
It  is  formed  as  a  black  precipitate  by  the  addition  of  ammonium 
or  sodium  sulfide  to  a  solution  of  a  ferrous  salt.  Ferrous  sulfide 
dissolves  easily  in  dilute  sulfuric  or  hydrochloric  acid,  the  reaction 
which  is  commonly  used  for  the  preparation  of  hydrogen  sulfide. 

*  Ferric  Sulfide,  Fe2S3,  forms  as  a  black  precipitate  on  adding 
ammonium  sulfide  or  sodium  sulfide  to  a  solution  of  a  ferric  salt : 

2  FeCl3  +  3(NH4)2S  =  Fe2S3  +  6  NH4C1 

Ferric  sulfide  reacts  with  an  ammoniacal  solution  of  zinc 
chloride  to  give  zinc  sulfide  and  ferric  hydroxide,  while  ferrous 
sulfide,  FeS,  reacts  scarcely  at  all  with  the  same  solution. 
(Stokes,  J.  Am.  Chem.  Soc.  29,  304.) 

Iron  Bisulfide,  or  Iron  Pyrites,  FeS2,  is  a  bright  yellow  mineral 
sometimes  called  "  fools'  gold  "  because  of  its  color  and  appear- 
ance. It  is  brittle  and  is  easily  recognized  by  its  burning  to 
sulfur  dioxide  and  ferric  oxide  when  heated  in  the  air.  It  is 
the  chief  source  of  sulfur  for  the  manufacture  of  sulfuric  acid. 

Ferric  Thiocyanate,  Fe(CNS)a.  When  ammonium  thio- 
cyanate,  NH4CNS,  is  added  to  a  solution  of  a  ferric  salt,  ferric 


^ COBALT  557 

thiocyanate  is  formed  and  imparts  a  deep  red  color  to  the  solu- 
tion. The  reaction  is  used  for  the  qualitative  detection  and 
sometimes  for  the  colorimetric  determination  of  iron. 

The  ferrocyanides  and  ferricyanides  have  been  described  in  a 
previous  chapter. 

Cobalt,  Co,  58.97.  Cobalt  is  usually  found  associated  with 
nickel  and  combined  with  arsenic  or  sulfur  or  both.  The  best 
known  mineral  is  smaltite  (CoNi)As2.  Cobalt  is  a  hard  white 
metal,  closely  resembling  iron.  It  is  malleable  and  ductile  and 
slightly  magnetic.  As  it  is  not  very  abundant  and  has  few 
properties  which  would  make  it  distinctly  more  valuable  than 
iron,  it  has  not  attained  any  considerable  commercial  use  as  a 
metal.  An  alloy  with  chromium  is  even  more  resistant  to  the 
attack  of  acids  than  nichrome  (p.  560)  and  offers  some  promise 
of  use  for  fruit  knives  and  for  spatulas  for  chemical  laboratories 
(Haynes,  J.  Ind.  and  Eng.  Chem.  2,  397).  Cobalt  melts  at  1478° 
and  has  a  specific  gravity  of  8.5. 

Compounds  of  Cobalt.  Oxides.  Cobalt  forms  cobaltous  and 
cobaltic  compounds,  corresponding  to  the  ferrous  and  ferric 
salts,  but  while  iron  tends,  on  the  whole,  to  pass  into  the  ferric 
state,  the  ordinary  compounds  of  cobalt  are  the  cobaltous  salts. 
The  four  oxides  are  cobaltous  oxide,  CoO,  cobaltic  oxide,  Co2O3, 
cobaltous-cobaltic  oxide,  CosO^  and  cobalt  dioxide,  CoO2. 

*  Cobaltous  Hydroxide,  Co(OH)2.     On  adding  potassium  hy- 
droxide to  a  solution  of  a  cobaltous  salt  a  blue,  basic  precipitate 
is  formed.     This  changes  on  boiling  to  pink  cobaltous  hydroxide, 
Co  (OH)  2.     If  cobaltous  hydroxide  is  heated  with  exclusion  of 
the  air,  it  changes  to  light  green  cobaltous  oxide,  CoO. 

*  Cobaltous  Chloride,  CoCl2.6  H2O,  dissolves  in  water  to  the 
reddish  or  pink  solutions  which  are  characteristic  of  all  cobaltous 
salts.     The  hydrated  salt  is  also  pink,  but  in  dry  air  or  on  warm- 
ing it  loses  its  water  of  hydration  and  changes  to  a  deep  green 
color.     This  is  made  the  basis  of  a  "  sympathetic  ink."     Draw- 
ings made  with  a  solution  of  cobalt  chloride  remain  nearly  in- 
visible in  moist  air,  but  come  out  to  a  clear  green  in  dry  air  or 
on  warming  gently. 


558  A  TEXTBOOK   OF  CHEMISTRY 

*  Cobalt  Sulfide,  CoS,  is  a  black  sulfide  which  dissolves  only 
slowly  in  dilute  hydrochloric  acid.     Nickel  sulfide,  NiS,  conducts 
itself  in  the  same  way,  and  the  property  is  often  made  the  basis 
of  a  partial  separation  of  these  metals  from  the  sul  fides  of  iron, 
zinc  and  manganese,  which  dissolve  easily  and  quickly  in  dilute 
acids.     The  difference  seems  to  be  due  rather  to  the  speed  of 
solution  than  to  a  marked  difference  in  the  solubilities  of  the 
sulfides  (A.  A.  Noyes,  Bray  and  Spear,  J.  Am.  Chem.  Soc.  30 
483).     Cobalt  sulfide  dissolves  easily  in  nitric  acid. 

*  Cobalt   Nitrate,    Co(NO3)2.6  H2O,    is    also    a    pink,   easily 
soluble  salt.     It  is  used  in  blowpipe  analysis.     Alumina  gives  a 
deep  blue  color  when  moistened  with  it  and  ignited,  while  zinc 
oxide  gives  a  green  color. 

Cobalt  Glass.  Cobalt  compounds  impart  an  intense  blue 
color  to  the  borax  bead,  and  this  is  used  as  a  delicate  test  for  the 
element.  They  give  a  similar  color  to  glass  and  are  used  in 
the  manufacture  of  blue  glass  and  in  decorating  porcelain.  Smalt 
is  a  deep  blue  silicate  of  potassium  and  cobalt  which  has  been 
disintegrated  by  pouring  the  melted  glass  into  water  and  after- 
ward ground  to  a  fine  powder.  It  has  a  variety  of  technical  uses. 

*  Potassium    Cobaltocyanide,   K4CoC6N6,    is    a    salt    closely 
analogous  to  potassium  ferrocyanide,  and  is  formed  in  a  similar 
manner  by  dissolving  cobaltous  cyanide,  CoC2N2,  in  a  solution 
of  potassium  cyanide. 

*  Potassium  Cobalticyanide,  K3CoC6N6,  is  formed  by  evaporat- 
ing a  solution  of  the  cobaltocyanide  exposed  to  the  air.     These 
compounds  correspond  to  the  potassium  ferro-  and  ferricyanides 
(p.  320).     Nickel  forms  no  similar  compounds 

Potassium  Cobaltinitrite,  K3Co(NO2)6.H2O,  forms  as  a  bright 
yellow,  difficultly  soluble  precipitate  on  adding  a  concentrated 
solution  of  cobaltous  acetate  and  sodium  nitrite,  or  an  adding  a 
solution  of  sodium  cobaltinitrite,  Na3Co(NO2)6,  to  a  solution  of 
a  potassium  salt.  The  reaction  is  used  for  the  detection  and 
quantitative  estimation  of  either  cobalt  or  potassium.  For  the 
determination  of  potassium  the  addition  of  silver  nitrate, 
AgNO3,  causes  the  formation  of  a  still  less  soluble  salt  in  which 


COBALT.    NICKEL.  559 


the    potassium   is   partly   replaced    by   silver,   ] 

(Burgess  and  Kamm,  J.  Am.  Chem.  Soc.  34,  652).     Nickel  forms 

no  similar  compounds. 

Cobalt  Ammines.  Cobalt  forms  a  very  great  number  of  com- 
plex compounds  with  ammonia.  The  compounds 

Luteocobalt  chloride,  [Co(NH3)6]  C13 

Purpureocobalt  chloride,     rCo(NH3)5~|  C\z  and 

L    ci 

Roseocobalt  chloride,         rCo(NH3)5l  C13 

L     H20     J 

may  be  given  as  illustrations.  In  these  compounds  the  cobalt 
combines  with  ammonia,  or  ammonia  and  water,  or  sometimes 
with  ammonia  and  chlorine  or  other  halogens  or  acid  radicals, 
to  form  complex  groups  which,  in  turn,  combine  with  acid 
radicals  to  form  salts.  Some  of  these  compounds  show  forms 
of  isomerism  recalling  that  of  the  hydrates  of  the  chromic 
chlorides  (p.  526).  A  study  of  these  compounds  by  Werner, 
especially,  has  led  to  the  development  of  new  views  of  valence 
which  are  important,  and  which  differ  considerably  from  the 
points  of  view  developed  from  the  study  of  organic  compounds. 
See  Werner,  Z.  anorg.  Chem.  3,  267-330  (1893) ;  and  Neuere 
Auschauungen  auf  dem  Gebiete  der  anorganischen  Chemie, 
1905. 

Nickel,  Ni,  58.68,  is  found  always  associated  with  some 
cobalt  and  combined  with  arsenic  or  sulfur,  or  both,  or  in  the 
form  of  a  silicate.  The  principal  sources  for  the  nickel  of  the 
world  have  been  Sudbury  in  Ontario  and  New  Caledonia,  an 
island  east  of  Australia.  At  Sudbury  the  ore  is  a  complex  sulfide 
containing  copper,  arsenic  and  other  metals.  The  metallurgical 
process  is  complicated  and  the  details  are  not  well  known  out- 
side of  the  works  where  they  are  carried  out.  The  first  step  is 
the  preparation  of  a  nickel-copper  matte  by  smelting  the  ore, 
somewhat  as  is  .done  with  copper  ores  (p.  429) . 

Nickel  is  a  white  metal  resembling  steel.  It  is  magnetic, 
melts  at  1452°  and  has  a  specific  gravity  of  8.8.  It  takes  a  high 


560  A  TEXTBOOK  OF  CHEMISTRY 

polish  and  does  not  rust  or  tarnish  so  readily  as  steel.  It  is 
extensively  used  for  plating  the  ornamental  parts  of  stoves, 
handle  bars  of  bicycles  and  many  other  articles.  For  this  pur- 
pose it  is  usually  deposited  from  a  solution  of  nickel  ammonium 
sulfate. 

Nickel  is  used  in  several  alloys,  the  most  important  being 
German  silver,  an  alloy  of  nickel,  copper  and  zinc,  which  is 
light  colored  and  used  as  a  basis  for  silver  plated  ware.  Ni- 
chrome,  an  alloy  of  nickel  and  chromium,  is  resistant  to  acids 
and  is  also  suitable  for  triangles  for  laboratory  use,  for  thermo- 
couples, and  for  electric  heating  devices.  Other  alloys  highly 
resistant  to  acids  have  been  prepared  and  there  is  a  prospect  of 
the  development  of  important  new  alloys  of  this  type.  The 
five-cent  coin  of  the  United  States  is  75  per  cent  copper  and 
25  per  cent  nickel. 

Compounds  of  Nickel.  Nickel  forms  three  oxides,  nickelous 
oxide,  NiO,  nickelic  oxide,  Ni2O3,  and  nickelous  nickelic  oxide, 
Ni3O4,  corresponding  to  the  similar  oxides  of  iron.  The  salts 
of  nickel  form  green  solutions  which  are  complementary  in  color 
to  the  salts  of  cobalt.  Mixtures  of  the  two  can  be  made  which 
are  nearly  colorless.  The  chloride  is  NiCl2-6  H^O  and  the 
sulfate  NiSO4.7  H2O. 

Nickel  Dimethylglyoxime, 

CH3—  C  =N—  O\       [~CH3—  C  =NOH 


CH3—  C  =N—  (X          CH3—  C  =NOH 

Nickel  forms  no  precipitate  with  potassium  nitrite,  and  cobalt 
can  be  separated  from  solutions  containing  it  by  this  means. 

CH3—  C=NOH 
It  is  precipitated  by  dimethylglyoxime,  ,  from 

CH3—  C=NOH 

an  ammoniacal  solution  or  from  a  solution  containing  a  weak 
acid,  in  the  form  of  a  scarlet-red,  highly  characteristic  compound, 
nickel  dimethylglyoximine.  Cobalt  forms  no  similar  precipitate, 
especially  if  the  ammoniacal  solution  is  first  shaken  in  the  air 
to  oxidize  the  cobalt  and  convert  it  into  a  complex  cobaltic 


NICKEL  561 

ammine.  The  formation  of  this  precipitate  furnishes  one  of 
the  best  means  for  the  detection  and  quantitative  estimation  of 
nickel.  The  precipitate  may  be  sublimed  without  decomposition 
by  careful  heating  (Chugaev,1  Ber.  38,  2520;  Z.  anorg.  Chem. 
46,  144;  Kraut,  Z.  angew.  Chemie,  19,  1793). 

Nickel  Carbonyl,  Ni(CO)4.  By  passing  carbon  monoxide 
over  finely  divided  nickel  at  a  temperature  below  80°,  nickel 
carbonyl  is  formed.  It  is  a  volatile  liquid  which  boils  at  43° 
and  whose  vapor  decomposes  explosively  at  60°.  Cobalt  forms 
no  similar  compound,  and  attempts  have  been  made  to  use  it 
industrially  for  the  separation  of  nickel  from  other  elements, 
but  these  attempts  have  not  met  with  much  success.  Iron 
forms  a  similar  compound,  iron  tetracarbonyl,  Fe(CO)4,  and  also 
a  pentacarbonyl,  Fe(CO)5,  but  these  are  much  less  stable  than 
the  nickel  carbonyl.  However,  it  seems  probable  that  traces 
of  these  iron  compounds  are  formed  from  the  carbon  monoxide 
of  illuminating  gas  and  are  the  cause  of  the  deposit  of  ferric 
oxide  sometimes  obtained  from  gas  burners. 

1  Also  written  Tschugaeff. 


CHAPTER  XXXIII 


THE   PLATINUM    METALS 

THE  following  table  gives  the  atomic  weights,  specific  gravity, 
melting  points  and  formulas  of  the  oxides  of  the  metals  of 
Group  VIII  of  the  Periodic  System : 

Fe                        Co  Ni 

Atomic  weight 55.84              58.97  58.68 

Specific  gravity 7.88                8.7  8.8 

Melting  point 1530°              1478°  1452° 

FeO               CoO  NiO 

n  .                                              ,  Fe304            Co3O4  Ni3O4 

'     '  Fe203             Co2O3  Ni203 
CoO2 

Ru                       Rh  Pd 

Atomic  weight 101.7              102.9  106.7 

Specific  gravity    .....          12.2                12.6  11.9 

Melting  point 2300°              1940°  1549° 

RuO              RhO  PdO 
,  Ru203           Rh203 

^Ru02            Rh02  Pd02 

RuO4 

Os                        Ir  Pt 

Atomic  weight 190.9              193.1  195.2 

Specific  gravity 22.48              22.4  21.4 

Melting  point       .....      2700°             2300°  1755° 

OsO 

Oxid                                             .Os203             Ir203  PtO 

1 ^Os02              Ir02  Pt02 

OsO4 

The  platinum  metals  are  found  almost  exclusively  in  the  free 
state,  alloyed  together  in  small  grains  or  nuggets.  Platinum 
forms  two  thirds  to  five  sixths  of  the  alloy. 

562 


THE  PLATINUM  METALS  563 

*  Ruthenium,  Ru,   101.7,  is   found   in   the   mineral   laurite, 
(RuOs)2S3,  as  well  as  in  the  natural  platinum  alloys  and  in 
osmium-iridium.     Ruthenium  monoxide,  RuO,  is  obtained  by 
heating  a  mixture  of  ruthenium  dichloride,  RuCl2,  and  sodium 
carbonate.     The  sesquioxide,  Ru2O3,  is  formed  when  the  metal 
is  heated  in  the  air.     It  will  be  seen  from  this  that  ruthenium 
resembles    iron    rather    than    platinum.     Ruthenium   dioxide, 
RuO2,  is  prepared  by  heating  ruthenium  in  a  current  of  oxygen. 
Ruthenium  tetroxide,  RuO4,  is  a  volatile  compound  formed  by 
fusing    ruthenium    with    potassium    hydroxide  and  potassium 
nitrate.     Its  odor  resembles  that  of  ozone.     Ruthenium  forms 
three  chlorides,   RuCl2,   RuCl3  and  RuCU,   and  such  double 
chlorides  as  K4RuCl6  (or  4  KCLRuCl2)  and  K2RuCl6.     The 
potassium  ruthenate,  K2RuO4.H2O,  which  gives  orange-red  solu- 
tions and  perruthenate,  KRuC>4.H2O,  giving  a  green  solution, 
correspond  to  the  manganates  and  permanganates. 

*  Rhodium,  Rh,  102.9,  is  much  less  easily  attacked  by  acids 
and  other  reagents  than  ruthenium.     It  is  hard  and  has  been 
sometimes  used  for  the  tips  of  gold  pens.     Rhodium  monoxide, 
RhO,  is  obtained  by  heating  the  hydroxide,  Rh  (OH)  3.     The 
sesquioxide,  Rh2O3,  is  prepared  by  heating  the  nitrate,  and  the 
dioxide,  RhO2,  by  fusing  rhodium  with  potassium  hydroxide 
and  nitrate.     The  chlorides  are  RhCl2  and  RhCl3.     The  double 
chloride  with  potassium  is  K2RhCl5.H2O  or  2  KCl.RhCl3.H2O. 
There  are  many  complex  salts,  such  as  roseorhodium  chloride, 

(~Rh(^TH')5lci3,  and  luteorhodium  chloride,  Rh(NH3)6Cl3. 
L        H2°  J 

Palladium,  Pd,  106.7,  is  always  present  in  the  natural  platinum 
alloys.  It  is  also  frequently  found  in  metallic  silver.  It  re- 
sembles platinum  or  silver  in  appearance  and  occupies  a  posi- 
tion somewhat  between  them  in  its  properties.  It  can  be  rolled 
into  foil  and  drawn  into  wire.  It  is  soluble  in  nitric  acid,  and 
finely  divided  palladium  will  dissolve  in  hydrochloric  acid. 

One  of  the  most  remarkable  properties  of  palladium  is  its 
absorption  of  hydrogen.  If  palladium  foil  is  heated  in  an  at- 
mosphere of  hydrogen  and  then  allowed  to  cool  in  a  current  of 


564  A  TEXTBOOK  OF  CHEMISTRY 

the  gas,  100  grams  of  the  metal  will  absorb  about  0.64  gram  of 
hydrogen  or  approximately  7  liters  of  the  gas.  As  the  volume 
of  100  grams  of  the  metal  is  only  8.4  cc.,  it  follows  that  the  metal 
absorbs  more  than  800  times  its  volume  of  the  gas.  This 
property  has  been  used  as  a  convenient  method  of  weighing 
hydrogen  for  the  determination  of  its  atomic  weight.  The 
hydrogen  absorbed  is  in  an  active  form.  It  is  oxidized  to 
water  at  once  in  contact  with  oxygen  or  the  air.  A  solution  of 
colloidal  palladium  also  has  a  similar  catalytic  effect  and  has 
been  used  with  hydrogen  for  the  reduction  of  organic  compounds. 
It  has  even  been  proposed  to  use  this  method  commercially  for 
the  reduction  of  liquid  or  semiliquid  fats  containing  glycerides  of 
unsaturated  acids  to  convert  these  into  solid  fats  which  are 
commercially  much  more  valuable. 

*  Palladium  forms  only  two  well-defined  oxides,  the  monoxide, 
PdO,  and  the  dioxide,  PdO2.     Palladium  dichloride,  PdCl2.2  H2O, 
is  obtained  by  dissolving  spongy  palladium  in  hydrochloric  acid. 
The  solution  is  reduced  by  hydrogen  and  is  sometimes  used  to 
absorb  that  gas  in  gas  analysis.     The  double  salt,  K2PdCl4, 
dissolves  in  water  to  a  dark  red  solution. 

Palladium  tetrachloride,  PdCl4,  is  known  only  in  solution 
and  is  not  very  stable.  The  double  salt,  K2PdCl6  or 
2  KCl.PdCl4,  is  difficultly  soluble  in  cold  water  and  crystallizes 
in  scarlet-red  octahedra.  Palladium  forms  a  series  of  ammines 
similar  to  those  of  cobalt  and  rhodium. 

*  Osmium,  Os,  190.9,  is  the  heaviest  substance  known.     The 
name  is  given  because  of  the  strong  odor  of  its  volatile  tetroxide, 
OsO4.     It  is  found  with  platinum  and  iridium  in  the  alloy  called 
osmium-iridium,  which  is  insoluble  in  aqua  regia.     From  this 
alloy  the  osmium  is  obtained  by  heating  in  a  current  of  oxygen, 
which  converts   the  osmium  into   volatile  osmium   tetroxide, 
OsO4.     The  reguline  metal  has  a  bluish  color  somewhat  resem- 
bling zinc. 

Osmium  gives  four  oxides,  OsO,  Os2O3,  OsO2  and  OsO4. 
The  tetroxide  is  a  white  solid  which  dissolves  slowly  in  water, 
but  volatilizes  from  the  solution.  The  vapor  has  a  disagree- 


THE  PLATINUM  METALS  565 

able,  chlorine-like  odor.  It  attacks  the  eyes  strongly  and  is 
extremely  poisonous.  It  melts  at  40°  and  boils  at  about  100°. 
It  is  often  called  an  acid,  but  has  no  acid  properties.  The 
aqueous  solution  is  sometimes  used  in  histology  to  stain  or  harden 
tissues. 

The  chlorides  of  osmium  are  OsCl2,  OsCl3  and  OsCU. 
The  double  salts  with  potassium  are  OsCl3.3  KC1.6  H2O,  or 
K3OsCl6.6  H2O  and  K2OsCl6.  Potassium  osmate,  K2OsO4.2  H2O, 
crystallizes  in  rose-red  or  violet  octahedra. 

*  Indium,  Ir,  193.1,  melts  about  550°  higher  than  platinum, 
and  as  it  does  not  oxidize  in  the  air  at  high  temperatures,  it  has 
proved  especially  useful  for  some  forms  of  chemical  apparatus. 
It  is  also  used  for  the  tips  of  gold  pens,  because  of  its  hardness. 
Its  color  is  between  those  of  silver  and  tin.     The  oxides  are  the 
sesquioxide,  Ir2O3,  and  the  dioxide,  IrO2.     Both  are  decomposed 
at  a  high  temperature  into  iridium  and  oxygen. 

The  chlorides  are  IrCl2,  IrCl3  and  .IrCl4.  The  double  salts 
are  K3IrCle.6  H2O  and  K2IrCle.  The  latter  forms  dark  red 
octahedra,  difficultly  soluble  in  water. 

Platinum,  Pt,  195.2,  is  very  much  the  most  important  of  the 
platinum  metals.  Its  very  high  melting  point  (1745°)  and  the 
fact  that  it  does  not  dissolve  in  nitric,  hydrochloric  or  sulfuric 
acid  make  it  an  almost  indispensable  metal  in  the  laboratory. 
The  fact  that  its  coefficient  of  expansion  is  almost  the  same  as 
that  of  some  kinds  of  glass  has  led  to  its  usje  for  the  leading-in 
wires  of  electric  light  bulbs.  Its  use  as  a  catalytic  agent  to 
cause  the  combination  of  hydrogen  and  oxygen,  and  also  the 
combination  of  sulfur  dioxide  and  oxygen  in  the  contact  process 
for  sulfuric  acid  (p.  175),  have  been  given.  Platinum  sponge  is 
used  as  the  filtering  material  in  the  Munroe-Neubauer  crucibles. 

*  Platinous  Chloride,  PtCl2,  is  a  greenish  compound  formed  by 
passing  chlorine  over  platinum  at  240°-250°.     It  dissolves  in 
hydrochloric  acid,  giving  chloroplatinous  acid,  H2PtCl4.     The 
potassium  salt,  K2PtCl4,  is  used  in  photography. 

Chloroplatinic  Acid,  H2PtCle,  is  usually  prepared  by  dissolving 
platinum  in  aqua  regia  and  evaporating  the  solution  repeatedly 


566  A  TEXTBOOK  OF  CHEMISTRY 

to  expel  the  excess  of  nitric  acid,  but  it  is  extremely  difficult  to 
obtain  a  pure  product  in  this  manner.  The  pure  compound  is 
best  prepared  by  dissolving  platinum  black  electrolytically  in 
hydrochloric  acid  (Weber,  J.  Am.  Chem.  Soc.  30,  29).  The 
acid  is  easily  soluble  in  water,  giving  a  yellow  or  reddish  yellow 
solution,  according  to 'the  concentration.  With  potassium  or 
ammonium  salts  the  solution  gives  yellow  precipitates  of  potas- 
sium chloroplatinate,  K2PtCle,  and  ammonium  chloroplatinate, 
(NH^PtCle-  These  compounds  are  much  used  in  analytical 
chemistry.  Similar  compounds,  many  of  which,  however,  are 
more  easily  soluble,  are  formed  with  organic  bases.  Silver 
nitrate,  AgNO3,  precipitates  silver  chloroplatinate,  Ag2PtCle, 
and  not  silver  chloride,  AgCl,  from  a  solution  of  chloroplatinic 
acid. 

Platinic  Chloride,  PtCU,  is  obtained  by  heating  chloroplatinic 
acid  in  a  current  of  chlorine  at  360°.  When  dissolved  in  water, 
it-  gives  the  compound  H2PtCl4O.4  H2O  or  PtCl4.5  H2O.  The 
first  formula  is  justified  by  the  fact  that  four  molecules  of  water 
can  be  readily  expelled,  but  the  fifth  cannot  be  removed  without 
loss  of  chlorine. 

*  Platinum  Bisulfide,  PtS2,  is  precipitated  on  passing  hydrogen 
sulfide  into  a  solution  of  chloroplatinic  acid.  It  is  a  black  pre- 
cipitate which  dissolves  in  ammonium  sulfide  as  ammonium 
sulfoplatinate. 

Platinum  forms  a  long  series  of  complex  ammines. 


INDEX 


INDEX 


Abscissas,  axis  of,  43. 

Absolute,  alcohol,  325 ;  potential  of 
elements,  436  ;  temperature,  39  ; 
units,  33;  zero,  40. 

Abstract  sciences,  4. 

Acetamide,  acid  in  liquid  ammonia, 
208. 

Acetanilide,  340. 

Acetic  acid,  acetyl  chloride  from, 
245  ;  manufacture,  properties,  con- 
stituent of  "  liquid  smoke,"  329  ; 
strength  illustrated,  386 ;  titration 
of,  389  ;  solubility  of  calcium  phos- 
phate in,  462 ;  insolubility  of 
calcium  oxalate  in,  465. 

Acetone,  formation,  preparation, 
uses,  328 ;  solvent  for  acetylene, 
293. 

Acetyl  chloride,  formed  from  acetic 
acid,  245. 

Acetylene,  endothermic,  292,  294; 
formation,  292,  preparation  from 
calcium  carbide,  293 ;  in  illu- 
minating gas,  295 ;  light  from, 
293  ;  liquid,  explosive,  294 ;  poly- 
merization, 294 ;  solution  in  ace- 
tone not  explosive,  uses,  294 ; 
tetrabromide,  293. 

Acheson,  graphite  in  colloidal  solu- 
tion, 277. 

Acid,  definition,  45,  168 ;  properties 
from  oxygen,  23 ;  strong,  defined, 
168. 

Acids  and  bases,  writing  equations 
for  reactions  between,  156 ;  degree 
of  ionization,  table,  383 ;  dibasic, 
defined,  183 ;  nomenclature  of, 
123;  organic,  structure,  328; 
strength  of,  definition,  167; 
strength  of,  illustration,  386; 
strength  of  in  relation  to  solu- 
bility of  sulfides,  168 ;  tribasic, 
defined,  183 ;  weak  and  strong, 
386. 

Acid  chlorides,   defined,    189. 

Acidimetry,  185. 

Acidity  or  alkalinity  of  indicators 
at  change  of  color,  table,  388. 

Actinium,  475 ;  series  of  elements, 
475. 

Adiabatic  cooling  of  air  at  higher 
levels,  232. 


Adjective  dyes,  342. 

Adsorption,  278. 

Affinity,  chemical,  29. 

Agate,  348. 

Air,  absorption  as  mixed  gas  by 
water,  228;]  a  mixture,  228; 
amount  of  fresh,  required  per 
hour,  231 ;  analysis  of  by  nitric 
oxide,  230 ;  calculated  weight 
of  1  liter,  229 ;  coefficient  of  ex- 
pansion, 38. 

Air,  composition  of  by  volume  and 
weight,  228 ;  demonstrated  by 
Lavoisier,  19 ;  determined  by 
hydrogen,  227 ;  determined  by 
mercury,  227 ;  determined  by 
phosphorus,  227. 

Air,  determination  of  moisture  in 
by  weighing,  by  dew  point,  and 
by  moist  bulb  of  thermometer, 
232;  liquefaction,  232;  puri- 
fication before  liquefaction,  234  ; 
sources  of  carbon  dioxide  in,  229 ; 
weight  of  gram  molecular  volume, 
228 ;  weight  of  1  liter,  229. 

Air-slaked  lime,  453. 

Alabaster,  457. 

Alberene,  349. 

Albite,   trisilicate,   356. 

Albumen,  343. 

Albumoses,  formed  in  digestion,  344. 

Alchemists,  name  for  silver,  444. 

Alcohol,  defined,  324 ;  manufacture, 
properties,  uses,  absolute,  dena- 
tured, 325. 

Alcoholic  beverages,  composition, 
325. 

Aldehydes,  327. 

Alfalfa,  fixation  of  nitrogen  by,  199. 

Alizarin,    manufacture,    341. 

Alkali  industry,  history  of,  400. 

Alkali  metals,  general  properties, 
395. 

Alkali-earth  metals,  general  proper- 
ties, 451. 

Alkalimetry,  185. 

Alkalinity  or  acidity  of  indicators 
at  change  of  color,  table,  388. 

Alkaloids,  342. 

Allotropic  forms,  definition,  98 ; 
of  phosphorus,  241 ;  of  sulfur, 
162. 


570 


INDEX 


Alloy  steels,  552. 

Alloys,  fusible,  269. 

Alum,  ammonium  gallium,  506 ; 
caesium,  rubidium,  424 ;  chrome, 
527. 

Alums,  500,  potassium,  500,  ammo- 
nium, ammonium  ferric,  chrome, 
rubidium,  501. 

Alumina,  blue  color  with  cobalt 
nitrate,  558. 

Aluminium  acetate,  mordant,  342 ; 
amalgam,  activity  of,  497  ;  bronze, 
from  electric  furnace,  495,  com- 
position, 497 ;  chloride,  prepara- 
tion, anhydrous,  hydrate,  use, 
498;  exercises,  507. 

Aluminium,  history  of  metallurgy 
of,  391 ;  occurrence,  formation 
of  shales,  clays,  soils,  494 ;  metal- 
lurgy, history,  495 ;  manufacture, 
496;  properties,  497;  alloys, 
thermite  process,  497  ;  compounds, 
498 ;  use  in  cast  iron,  543  ;  use  in 
metallurgy,  391. 

Aluminium  hydroxide,  precipitation, 
499  ;  preparation  from  clay,  496  ; 
base  and  acid,  499 ;  fluoride,  prep- 
aration, 499 ;  metachlproanti- 
monate,  268 ;  oxide,  dissolved 
by  sodium  pyrosulfate,  408  ;  oxide, 
preparation  for  manufacture  of 
aluminium,  495,  496 ;  oxide,  oc- 
currence, 494 ;  artificial,  499 ; 
uses,  500 ;  solution  in  sodium 
pyrosulfate,  500  ;  sulfate,  hydroly- 
sis, preparation,  use,  499,  500. 

Amalgam,  ammonium,  420 ;  sodium, 
487. 

Amalgamated  zinc,  conduct  toward 
acids,  481. 

Amalgamation  process  for  silver, 
441. 

Amalgams,  486. 

Amethyst,  348. 

Amide,  definition  of,  206 ;  ions  in 
liquid  ammonia,  208. 

Amine,  definition  of,  properties,  205. 

Amino  acids,  formed  in  digestion, 
344. 

Ammines,  cobalt,  559. 

Ammonia,  "  associated  "  liquid,  204  ; 
combination  with  acids,  202  ;  deriv- 
atives of,  205 ;  detection  with 
Nessler's  reagent,  492. 

Ammonia,  determination  of  com- 
position by  volume  by  action  of 
chlorine  on,  208  ;  by  decomposition 
and  recombination  with  electric 
discharge,  209. 

Ammonia,  deviation  from  Boyle's 
law,  35 ;  formation,  201 ;  formed 


by  action  of  zinc  on  nitric  acid, 
213  ;  liquid,  solutions  in,  207. 

Ammonia,  preparation  by  hydrolysis 
of  a  nitride,  202  ;  from  ammonium 
sulfate,  202  ;  from  aqua  ammonia, 
202. 

Ammonia,  properties,  solubility,  202 ; 
reaction  between  chlorine  and,  209  ; 
synthesis  of,  201. 

Ammonia  soda  process,  discovery, 
400. 

Ammoniacal  gas  liquors,  202. 

Ammonio-cadmium  sulfate,  492;  -cu- 
pric  sulfate,  434  ;  -cuprous  chloride, 
433  ;  -zinc  sulfate,  492. 

Ammonium,  in  ammonium  amal- 
gam, 420 ;  bicarbonate,  423  ;  bi- 
carbonate, use  in  ammonia  soda 
process,  412 ;  carbonate,  com- 
mercial, preparation,  composition, 
use,  423 ;  carbonate,  formation, 
hydrolysis,  use,  423  ;  chloroaurate, 
450. 

Ammonium  chloride,  formation  from 
ammonia  and  hydrochloric  acid, 
203 ;  manufacture,  dissociation, 
volatilization  of  dry  without  dis- 
sociation, relation  to  Avogadro's 
hypothesis,  421 ;  recovery  of  am- 
monia from,  413. 

Ammonium  chloroplatinate,  566,  423. 

Ammonium  chloroplumbate,  518  ; 
citrate,  use  in  analysis  of  fertilizers, 
331 ;  citrate,  use  in  determining 
citrate-soluble  phosphoric  acid, 
461 ;  ferric  citrate,  use  with  potas- 
sium ferricyanide  in  blue  prints, 
331. 

Ammonium  hydroxide,  dissociation  to 
ammonia  and  water,  204 ;  equili- 
bria in  solutions  of,  420 ;  ioniza- 
tion,  203 ;  structure  according  to 
election  theory,  207 ;  titration  of, 
389. 

Ammonium  hydrosulfide,  prepara- 
tion, formation  of  polysulfides 
from,  421 ;  magnesium  phosphate, 
decomposition  of,  252  ;  molybdate, 
use  to  determine  phosphorus,  529, 
530. 

Ammonium  nitrate,  decomposition  of 
exothermic,  use  in  explosives,  215, 
423  ;  preparation  of  nitrous  oxide 
from,  214 ;  preparation,  proper- 
ties, use,  422. 

Ammonium  nitrite,  properties,  423 ; 
oxalate,  use  to  precipitate  cal- 
cium, 330 ;  phosphomolybdate, 
use  to  determine  phosphoric  acid, 
529,  530 ;  polysulfides,  formation, 
use,  422. 


INDEX 


571 


Ammonium  salts,  theory  of,  203; 
sodium  hydrogen  phosphate,  use, 
423  ;  sulfarsenite,  261. 

Ammonium  sulfate  from  ammoniacal 
gas  liquors,  202  ;  preparation,  use, 
422  ;  use  as  fertilizer,  199. 

Ammonium  sulfide,  preparation,  421, 
sulfide,  formation  of  polysulfides 
from,  uses,  422 ;  sulfostannate, 
512;  trinitride,  structure,  221. 

Ammono-dimercuric  iodide,  use  in 
testing  for  ammonia,  493. 

Ammono-mercuric  chloride,  492  ; 
compounds,  492  ;  nitrate,  492. 

Amorphous,  definition,  162. 

Ampere  defined,  33. 

Amphibole,  metasilicate,  355. 

Amphoteric  compounds,  definition  of, 
206,  483. 

Amyl  acetate,  use  in  lacquers,  338. 

Analysis,  definition,  66 ;  hydrogen 
sulfide  basis  of  groups  in,  166. 

Analytical  chemistry,  groups  of,  166. 

Andrews,  critical  temperature,  232. 

Anesthesia,  produced  by  nitrous 
oxide,  215. 

Anhydrite,  457,  458,  soluble,  458 ; 
conditions  for  formation  of,  458, 
460 ;  vapor  pressure  of  systems  con- 
taining, 459. 

Aniline,  preparation,  uses,  340. 

Animal  charcoal,  278. 

Animal  foods,  347. 

Anion,  definition,  48. 

Anode,  definition,  47,  113;  deposit 
of  silver  peroxide  on,  443. 

Anthracene,  alizarin  from,  341 ;  from 
coal  tar,  use,  295. 

Anthracite  coal,  composition,  280. 

Antifebrine,  340. 

Antifriction  metals,  264. 

Antimonic  acids,  267. 

Antimonious  acid,  265. 

Antimony,  chlorides  of,  267 ;  hy- 
droxide, 265  ;  oxides  of,  265. 

Antimony,  occurrence,  preparation, 
263;  properties,  uses,  alloys,  ex- 
plosive, 264. 

Antimony  oxychloride,  267;  penta- 
chloride,  267;  pentoxide,  265; 
pentasulfide,  268. 

Antimony  tetrachloride,  267;  tetra- 
chloride, endothermic,  267 ;  tet- 
roxide,  265 ;  trichloride,  explosive 
antimony  from,  264 ;  trichloride, 
preparation,  properties,  hydrolysis, 
267  ;  trioxide,  265 ;  trioxide,  tartar 
emetic  from,  266 ;  trisulfide,  268. 

Antimonyl,  266;  chloride,  267;  po- 
tassium tartrate,  266 ;  sulfate,  266. 

Antipyrine,  340. 


Antitoxins,  345. 

Apatite,  153,  452. 

Apollinaris  water,  309. 

Aqua  ammonia,  202,  203. 

Aqua  regia,  213  ;  use  to  oxidize  sulfur 
of  sulfides,  213. 

Aragonite,  452. 

Argentum,  11. 

Argol,  330. 

Argon,  atomic  weight,  236 ;  coeffi- 
cient of  expansion,  38 ;  discovery, 
235 ;  molecular  weight,  236 ;  per 
cent  in  ah*,  228 ;  properties,  236. 

Argyrodite,  discovery  of  germanium 
in,  361. 

Arsenic  acid,  oxidizing  agent,  259 ; 
preparation,  properties,  salts,  259  ; 
transformation  in  steps  to  arsenic 
pentasulfide,  261. 

Arsenic  disulfide,  260. 

Arsenic  from  smelting  copper  ores, 
256;  occurrence,  256. 

Arsenic  pentasulfide,  260  ;  pentoxide, 
259  ;  "  poison  "  to  platinum  catal- 
ysis for  preparation  of  sulfur 
trioxide,  175 ;  preparation,  prop- 
erties, uses,  257  ;  sulfides  of,  260. 

Arsenic  trioxide,  conversion  to  col- 
loidal arsenic  trisulfide,  261 ;  tri- 
oxide, formation,  properties,  258 ; 
trisulfide,  260;  trisulfide,  col- 
loidal, 261. 

Arsenious  acid,  salts,  259 ;  oxide, 
oxidation  by  nitrogen  trichloride, 
224 ;  oxide,  standard  in  iodimetry, 
260. 

Arsenites,  259. 

Arsenopyrite,  256  ;  arsenic  from,  257. 

Arsine,  Marsh's  test,  257 ;  com- 
pared with  ammonia,  243. 

Asbestos,  diaphragm  of  for  alkali 
manufacture,  401 ;  metasilicate, 
355 ;  platinized,  preparation  of,  62. 

Assaying,  440. 

Association,  water  and  ammonia,  204. 

-ate,  suffix,  use,  47  ;  use  for  salts,  124. 

Atmosphere,  exercises,  239. 

Atomic  theory,  14;  volumes,  curve 
for,  137 ;  volumes,  relation  to 
periodic  system,  136 ;  weight  of 
chlorine,  determination  of,  130. 

Atomic  weights,  selected  by  law  of 
Dulong  and  Petit,  397 ;  selection 
of,  16,  92 ;  table  of,  10 ;  unit  for, 
68. 

Atoms,  probably  complex  aggre- 
gates, 138 ;  structure  of,  473. 

Atropa  belladonna,  atropine  from, 
343. 

Atropine,  343. 

At  water,  respiration  calorimeter,  313. 


572 


INDEX 


-Auric  acid,  449. 

Aurum,  11. 

Austenite,  relation  to  tempering  of 

steel,  546. 
Avogadro's  law,  89,  91 ;  exercises,  99  ; 

related    to    law    of    Dulong    and 

Petit,    397 ;     relation    to    laws    of 

Boyle  and  Charles,  94. 
Az-,  prefix  derived  from  azote,  220. 
Azo-,  prefix  derived  from  azote,  220. 
Azoimide,  see  hydronitric  acid,  223. 
Azote,  name  for  nitrogen,  220. 

Babbitt  metal,  269. 

Bacteria,  killed  by  radiations  from 
radioactive  elements,  476 ;  re- 
moval from  water,  83 ;  nitrifying, 
199. 

Badische  Anilin  Soda  Fabrik,  syn- 
thetic ammonia,  201 ;  manufac- 
ture of  indigo,  341. 

Baeyer,  synthesis  of  indigo,  341. 

Baker,  atomic  weightof  tellurium,  190. 

Baking  powder,  acid  potassium  tar- 
trate  in,  330. 

Baking  soda,  412. 

Banca  tin,  508. 

Barite,  162,  468 ;  use,  470. 

Barium  carbonate,  dissociation  pres- 
sure, manufacture  of  barium  oxide 
from,  468;  chloride,  470;  chro- 
mate,  528;  exercises,  476;  flame 
color,  471. 

Barium,  hydroxide,  properties,  uses, 
470  ;  nitrate,  preparation,  use,  470  ; 
nitrate,  barium  oxide  from,  469 ; 
occurrence,  compounds,  468  ;  oxide, 
manufacture,  468  ;  uses,  469. 

Barium  peroxide,  contrast  with  lead 
dioxide,  518 ;  dissociation  pres- 
sure, use  to  prepare  oxygen,  for 
hydrogen  peroxide,  469  ;  peroxide 
hydrate,  469 ;  hydrogen  peroxide 
from,  84. 

Barium  silicofluoride,  insoluble,  350 ; 
sulfate,  use,  constituent  of  litho- 
pone,  470  ;  sulfate,  solubility,  471 ; 
sulfide,  preparation,  use,  470. 

Barometer,  correction  of  readings  for 
altitude,  37  ;  correction  of  readings 
for  glass  and  brass  scales,  36; 
correction  of  readings  for  latitude, 
37. 

Bases,  definition,  121 ;  derived  from 
ammonia,  339 ;  how  formed  from 
ammonia  and  amines,  206 ;  weak 
and  strong,  386 ;  and  acids,  writing 
equations  for  reactions  between, 
156. 

Basicity,  defined,  183. 

Batteries  electric,  use  of  zinc  in,  481. 


Battery  galvanic,  reverse  of  electro- 
lytic cell,  439. 

Bausfield,  value  of  the  calorie  at 
different  temperatures,  33. 

Bauxite,  494. 

Baxter,  separation  of  praseodymium, 
and  neodymium,  504. 

Beads,  borax  and  sodium  metaphos- 
phate  with  blowpipe,  304. 

Bearings  in  machinery,  phosphor 
bronze  for,  431. 

Becquerel,  discovery  of  rays,  471. 

"  Bee  hive,"  coke  ovens,  278. 

Beef,  use  of  potassium  nitrate  in 
salt,  418. 

Beet  sugar,  333. 

Bell  metal,  509. 

Benedict,  respiration  calorimeter,  313. 

Benzaldehyde,  328. 

Benzene  in  illuminating  gas,  295 ; 
from  coal  tar,  294  ;  properties,  295  ; 
structure,  285. 

Benzine,  289. 

Benzoic  acid,  cocaine  a  derivative  of, 
343  ;  occurrence,  manufacture,  use 
as  a  food  preservative,  331. 

Beryl,  451. 

Beryl,  metasilicate,  355. 

Beryllium  carbonate,  452 ;  chloride, 
451;  hydroxide,  451 ;  nitrate-,  452  ; 
occurrence,  properties,  compounds, 
451 ;  sulfate,  451. 

Berzelius,  determination  of  composi- 
tion of  water,  69  ;  experience  with 
hydrogen  selenide,  190 ;  use  of 
isomer,  511. 

Bessemer  converter,  547 ;  use  in 
metallurgy  of  copper,  429. 

Bessemer  steel,  history,  547  ;  acid  and 
basic  or  Thomas-Gilchrist,  548; 
soaking  pits,  548. 

Bi-bivalent  salts,  law  of  solubility 
product  not  general  for,  378. 

Bicarbonate  ion,  ionization  of,  310. 

Bicarbonates,  formation  from  car- 
bonic acid,  310.  • 

Bimolecular  reactions,  150. 

Binary  compounds,  nomenclature  of, 
29. 

Biological  sciences,  4. 

Biscuit,  forms  for  earthenware,  501. 

Bisdiazoacetic  acid,  use  in  preparing 
hydrazine,  222. 

Bismuth,  alloys  of,  269  ;  basic  nitrates 
of,  270  ;  in  crude  copper,  430  ;  melt- 
ing point  lowered  by  pressure,  269. 

Bismuth,  occurrence,  properties,  uses, 
268;  nitrate,  preparation,  hy- 
drolysis, 270  ;  oxides  of,  269  ;  oxy- 
chloride,  270;  "  subnitrate,"  270; 
trichloride,  269  ;  trisulfide,  270. 


INDEX 


573 


Bismuthyl  chloride,  270. 

Bituminous  coal,  calculation  of  heat 
of  combustion  of,  44  ;  composition, 
280. 

Bivalent,  definition,  64. 

Black  ash  in  Leblanc  soda  process, 
411. 

Black  oxide  of  manganese,  history  of 
uses,  535. 

Blast,  heating  for  blast  furnace,  542 ; 
dry,  542. 

Blast  furnace,  541 ;  gas,  298 ;  per- 
centage composition,  299 ;  slag, 
use  for  cement,  454. 

Bleaching  powder,  124 ;  manufac- 
ture, 455  ;  properties,  uses,  456  ; 
use  in  purifying  water,  83. 

Bleaching  by  sulfur  dioxide,  173 ; 
with  chlorine,  106. 

Blindness  caused  by  methyl  alcohol, 
325. 

Blowpipe,  construction  and  use,  304 ; 
oxyhydrogen,  61. 

Blueing,  321. 

Blue-print  paper,  331. 

Blue  vitriol,  433. 

Bodenstein,  decomposition  of  hydrio- 
dic  acid,  148  ;  heat  of  formation  of 
hydriodic  acid,  153. 

Body,  definition,  7. 

Bohemian  glass,  467. 

Boiling  point,  criterion  of  pure  sub- 
stance, 12 ;  of  solutions  and  os- 
motic pressure,  360. 

Boisbaudran,  discovery  of  samarium, 
505. 

Bolivia,  tin  from,  508. 

Bomb  calorimeter,  25. 

Bone  ash,  241. 

Bone  black,  278. 

Borax  beads  in  blowpipe  flame,  304 ; 
use  in  blowpiping,  366. 

Borax,  occurrence  in  California,  365 ; 
glass,  366 ;  uses,  367. 

Boric  acid,  in  Tuscany,  365 ;  prepa- 
ration, properties,  uses,  366 ; 
use  as  food  preservative,  366 ;  tests 
for,  367,  368. 

Bornite,  428. 

Boron,  exercises,  368 ;  fluoride, 
hydrolysis,  367;  nitride,  367;  oc- 
currence, preparation,  properties, 
365  ;  sulfide,  367  ;  trichloride,  367  ; 
trioxide,  365. 

Borosilicate  glasses,  467. 

Boyle,  law  of,  34. 

Boyle's  law,  deviation  of  gases  from, 
35  ;  illustration  of,  35. 

Brandt,  discovery  of  phosphorus,  242. 

Brandy,  325. 

Brass,  431. 


Brazil,  diamonds  from,  275. 

Breath,    source    of    carbon    dioxide, 

229. 

Brick,  501. 
Brimstone,  roll,  161. 
Brines,  evaporation  of,  405. 
Britannia  metal,  264,  509. 
British  thermal  unit  (B.  T.  U.),  26. 
Bromate,  sodium,  144. 
Bromine,  occurrence,  140 ;    origin  of 

name,      141 ;      preparation,      140 ; 

properties,  141 ;   uses,  142. 
Bronze  age,  relation  to  metallurgy, 

390;    phosphor,  431. 
Bronzes,  431,  509  ;  sodium  tungstate, 

531. 

Brownian  movement,  96. 
Bullion,    lead,    separation    of    silver 

from,  440. 
Bunsen,  discovery  of  rubidium  and 

caesium,   424 ;    discovery   of   spec- 
trum analysis,  424. 
Bunsen  burner,  nature  of  flame,  300  ; 

burner,  temperature  of  flame,  303  ; 

flame,  separated,  301. 
Burgess,  table  of  melting  points  of 

elements,  373. 
Burgess,    use    of    potassium    silver 

cobaltinitrite  in  analysis,  559. 
Burns  produced  by  radioactive  ele- 
ments, 476. 

Butadiene,  structure,  285. 
Butane,  normal,  structure,  284. 
Butene,  283. 
Butine,  283. 

By-products  coke  ovens,  279. 
By-products,  importance  in   Leblanc 

soda  process,  401. 

Cadmium,  abnormal  ionization  of,  491 ; 
compounds,  effect  of  ammonium 
hydroxide  .on  solutions  of,  491 ; 
hydroxide,  484. 

Cadmium,  occurrence,  properties, 
483,  use  in  fusible  alloys,  484 ; 
separation  from  copper  with  po- 
tassium cyanide,  435  ;  sulfate,  484  ; 
sulfate,  use  in  Weston  cell,  438; 
sulfide,  conduct  toward  acids  and 
potassium  cyanide,  484 ;  sulfide, 
solubility,  491. 

Caesium  alum,  424 ;  chloroiodide, 
use  in  purification  of  caesium,  424. 

Caesium  chloroplatinate,  393,  424. 

Caesium,  discovery,  properties,  424 : 
tetrachloroantimonate,  267. 

Cailletet,  liquefaction  of  air,  233. 

Calamine,  481. 

Calamine,  orthosilicate,  355. 

Calcite,  452. 

Calcium  acetate,    manufacture,  use, 


574 


INDEX 


465;  bicarbonate,  hard  waters, 
310. 

Calcium  carbide,  hydrolysis  to  cal- 
cium hydroxide  and  acetylene, 
293  ;  manufacture,  hydrolysis,  use, 
462 ;  calcium  cyanamide  from,  462. 

Calcium  carbonate,  dissociation,  phase 
rule,  453  ;  dissociation  pressures  of, 
453  ;  in  mortar,  454  ;  solubility  in 
pure  water  and  in  water  containing 
carbonic  acid,  hard  waters,  463. 

Calcium  chlorate,  formation,  use  in 
making  potassium  chlorate,  456 ; 
chlorate,  use  to  prepare  potassium 
chlorate,  127 ;  chloroaurate,  450. 

Calcium  chloride,  by-product  in  am- 
monia soda  process,  413 ;  effect 
on  the  formation  of  gypsum,  460 ; 
moisture  left  in  gas  by,  54  ;  prep- 
aration, hydrates,  by-product  in 
ammonia  soda  process,  as  drying 
agent,  uses,  455. 

Calcium  cyanamide,  manufacture, 
hydrolysis,  462;  use,  463;  exer- 
cises, 476 ;  flame  color,  471 ; 
fluoride,  properties,  uses,  456 ; 
fluoride,  occurrence,  153  ;  hydride, 
452 ;  hydrosulfide,  action  of  car- 
bonic acid  on,  457 ;  hydroxide, 
in  setting  of  cement,  455 ;  hy- 
pochlorite,  manufacture,  455,  prop- 
erties, uses,  456 ;  manganite, 
formation  in  Weldon  process,  103  ; 
metachloroantimonate,  268. 

Calcium  nitrate,  formation  in  soil, 
199  ;  occurrence,  potassium  nitrate 
from,  418 ;  preparation,  manu- 
facture, uses,  460. 

Calcium  nitride,  formation,  hydroly- 
sis, 452 ;  occurrence,  preparation, 
properties,  452 ;  orthoplumbate, 
516. 

Calcium  oxalate,  precipitation,  solu- 
bility in  strong  acids,  insolubility 
in  weak  acids,  use  as  test  for 
calcium  and  oxalic  acid,  465. 

Calcium  oxide,  manufacture,  452; 
equilibrium  with  carbon  dioxide 
and  calcium  carbonate,  453. 

Calcium  phosphates,  occurrence,  460 ; 
superphosphates,  461 ;  solubility, 
relation  to  fertilizing  value,  461, 
462. 

Calcium  phosphate,  solubility  in 
weak  acids,  462;  selection  of 
atomic  weight,  397  ;  silicate,  prep- 
aration, occurrence,  466;  stearate, 
332. 

Calcium  sulfate,  forms,  uses,  457; 
hydrates  of  and  the  phase  rule, 
458;  in  cement,  454;  in  hard 


waters,  463 ;  permanent  hardness 
from,  311. 

Calcium,  sulfite,  acid,  preparation, 
use  in  paper  manufacture,  457 ; 
sulfite,  acid,  use  in  paper  making, 
175;  sulfida  in  Leblanc  soda  pro- 
cess, 411 ;  sulfide,  preparation, 
hydrolysis,  456,  recovery  of  sulfur 
from  by  Chance  process,  457 ;  super- 
phosphate, manufacture,  use,  anal- 
ysis, 461. 

Calculation  of  formula  of  mineral, 
356 ;  of  relative  speed  of  reactions 
at  equilibrium,  151. 

Calomel,  preparation,  vapor  density, 
molecular  weight,  uses,  489. 

Calorie,  definition,  26 ;  value  at 
different  temperatures,  33 ;  value 
in  joules,  33. 

Calorimeter,  25. 

Calorimeter,  respiration,  313,  muscu- 
lar energy  in,  315;  mental  work  in, 
316;  cuts  of,  314,  315. 

Camphor,  in  celluloid,  338. 

Candle  power,  standard,  293,  of 
illuminating  gas,  296,  of  acetylene, 
293,  of  Welsbach  light,  296. 

Cane  sugar,  333. 

Cannel  coals,  281. 

Canon  Diablo,  diamonds  in  meteo- 
rite from,  274. 

Caramel,  333. 

Carat,  defined,  448. 

Carbohydrates,  defined,  332. 

Carbolic  acid,  see  phenol. 

Carbon,  amorphous,  heat  of  com- 
bustion, 276 ;  preparation,  prop- 
erties, 277. 

Carbon  bisulfide  and  nitric  oxide, 
flame  of,  217 ;  preparation,  prop- 
erties, uses,  formation  of  sulfo- 
carbonates  from,  317. 

Carbon,  chemical  properties  of,  281 ; 
cycle  of  in  nature,  312. 

Carbon  dioxide,  amount  formed  in 
world  by  burning  coal,  229 ; 
amount  in  the  ocean,  230  ;  amount 
in  air  kept  constant  by  ocean, 
230 ;  coefficient  of  expansion,  38 ; 
conditions  for  escape  from  solu- 
tions, 376. 

Carbon  dioxide,  critical  temperature 
of,  233;  density,  308;  accumula- 
tion in  wells  and  caves,  309, 
diffusion  in  rooms,  309  ;  deviation 
from  Boyle's  law,  35 ;  exercises, 
322 ;  formerly  considered  poison- 
ous, 231 ;  from  burning  charcoal,  23. 

Carbon  dioxide,  limit  for  in  ventilated 
rooms,  231 ;  per  cent  in  air,  228 ; 
preparation  from  calcium  car- 


INDEX 


575 


bonate,  sodium  bicarbonate,  mag- 
nesium carbonate,  306 ;  proper- 
ties, isothermals,  307. 

Carbon  dioxide,  ratio  of  specific 
heats  of,  237  ;  reduction  in  plants, 
312 ;  removed  from  air  by  plants, 
229  ;  solubility,  309  ;  sources,  313  ; 
sources  of  in  air,  229. 

Carbon,  effect  on  decomposition  of 
barium  carbonate,  468 ;  electrodes, 
279 ;  gas,  279 ;  heat  of  combus- 
tion, 27  ;  in  steel,  use  of  potassium 
cupric  chloride  in  determining,  432. 

Carbon  monoxide,  absorbed  by  cu- 
prous chloride,  433  ;  coefficient  of 
expansion,  38 ;  deviation  from 
Boyle's  law,  35 ;  effect  of  forma- 
tion on  decomposition  of  barium 
carbonate,  468 ;  formation  from 
calcium  oxalate,  466. 

Carbon  monoxide,  formation  in 
burning  coal,  preparation  from 
oxalic  acid,  311 ;  properties,  failure 
to  burn  when  dry,  311 ;  in  water 
gas,  danger,  297;  poisonous,  312; 
formation  of  sodium  formate  from, 
312. 

Carbon,  occurrence,  number  of  com- 
pounds, importance,  273 ;  oxy- 
chloride,  316 ;  oxysulfide,  forma- 
tion from  thiocyanates,  properties, 
hydrolysis,  318 ;  slowness  of  reac- 
tion, 280;  suboxide,  316;  tetra- 
chloride  from  methane,  287 ;  va- 
lence, 282. 

Carbonate,  ammonium,  423. 

Carbonate  ion,  r61e  in  decomposition 
of  carbonates,  375. 

Carbonates,  decomposition  by  acids, 
theory,  375 ;  formation  from  car- 
bonic acid,  310 ;  oxides  prepared 
from,  392. 

Carbonic  acid,  carbonates  and  bi- 
carbonates  from,  310 ;  deter- 
mination of  free  and  combined, 
464 ;  formation,  properties,  309. 

Carbonyl  chloride,  preparation,  prop- 
erties, 316,  hydrolysis,  urea  from, 
317. 

Carborundum,  349. 

Carboxyl,  characteristic  group  of 
organic  acids,  328. 

Carnallite,  414,  478;  rubidium  in, 
424. 

Carnotite,  uranium  in,  531. 

Caro's  acid  (permonosulfuric  acid), 
188. 

Casein,  a  protein,  343. 

Cassiterite,  508. 

Cast  iron,  composition,  gray,  white, 
chilled,  543  ;  analyses,  544. 


Castner-Kellner  apparatus  for  alkali 

manufacture,  402. 
Cast  steel,  545. 
Catalysis,    28;     definition,    62;     in 

sulfuric  acid  manufacture  by  oxides 

of  nitrogen,    178 ;    of  preparation 

of  sulfur  trioxide,  175 ;  of  synthesis 

of  ammonia,  201. 
Catalyzer,  copper  chloride  for  Deacon 

process,  103. 

Cathode,  definition,  47,  113. 
Cation,  definition,  48. 
Cavendish,  analysis  of  air  by  nitric 

oxide,    230 ;       nearly     discovered 

argon,  235. 
Celestite,  468. 
Celluloid,  338. 
Cellulose,   use  as  fuel,   as  food,   in 

paper,  337. 
Cement,    dental,    composition,   482 ; 

manufacture,     composition,     454; 

setting,  455. 
Cementation  steel,  545. 
Cementite,  iron  carbide,  546. 
Centimeter-gram-second  system,  33. 
Cerargyrite,  439. 
Cereal,  composition  of,  8. 
Cerium,  alloy  with  iron,  364 ;  group 

of   rare   earths,   503 ;     occurrence, 

properties,    362 ;     oxides,    sulfate, 

double  sulfate  with  sodium,  364; 

phosphate,  363. 
Chalcedony,  348. 
Chalcocite,  428. 
Chalcopyrite,  428. 

Chamberlain,  use  of  copper  to  pre- 
vent corrosion  of  iron,  550. 
"  Chamber  process  "  for  sulfuric  acid, 

177. 

Chameleon  solution,  536. 
Chance  process  for  recovering  sulfur, 

457. 
Chapin,  separation  of  praseodymium 

and  neodymium,  504. 
Charcoal,   animal,   278 ;    burning  in 

oxygen,    23 ;     composition,    280 ; 

manufacture,  277 ;  properties,  uses, 

adsorption  by,  278. 
Charles,  law  of,  38. 
Chemical  action  of  radioactive  rays, 

475. 

Chemical  activity  in  solutions,  81. 
Chemical    affinity,    29 ;    relation    to 

speed  of  chemical  reactions,  149. 
Chemical    energy,    nature    of,    27; 

defined,  34. 
Chemical  reactions,   equilibrium  in, 

108 ;     effect   of    concentration   on, 

24 ;     speed   of,    149 ;   unimolecular 

and  bimolecular,  150. 
Chemistry,  definition,  5 ;  study  of,  18. 


576 


INDEX 


Chicago,  typhoid  fever  in,  from 
water,  83. 

Chili  saltpeter,  210 ;  potassium  ni- 
trate from,  418. 

Chlorates,  127. 

Chloric  acid,  127 ;   structure,  130. 

Chloride,  acid,  hydrolysis  of,  317 ; 
of  lime,  manufacture,  455 ;  proper- 
ties, uses,  456. 

Chlorides,  hydrolysis  of,  115;  of 
acids,  189 ;  preparation  by  use  of 
sulfur  monochloride,  188. 

Chlorination  process  for  gold,  446. 

Chlorine  and  oxygen,  comparison  of 
heats  of  combination,  108 ;  action 
on  ammonia,  209 ;  bleaching  by, 
106;  burning  in  hydrogen,  118. 

Chlorine,  combination  with  other  ele- 
ments, 104 ;  determination  of  atomic 
weight  of,  130;  effect  of  light  on 
reaction  with  hydrogen,  105 ;  effect 
of  moisture  on  combination  with 
other  elements,  105  ;  exercises,  116 ; 
from  electrolysis  in  alkali  manufac- 
ture, 402. 

Chlorine  hydrate,  phases,  107;  list 
of  oxides  and  oxygen  acids  of,  123  ; 
occurrence,  100 ;  dioxide,  127. 

Chlorine,  preparation  by  electrolysis 
of  sodium  chloride,  100 ;  by  the 
Deacon  process,  102 ;  from  hydro- 
chloric acid  and  manganese  dioxide, 
101 ;  from  hydrochloric  acid  and 
potassium  permanganate,  102  ;  by 
oxidation  of  hydrochloric  acid,  100 ; 
by  Weldon  process,  102. 

Chlorine,  properties,  104;  reaction 
with  water,  106. 

Chlorites,  127. 

Chloroaurates,  450. 

Chloroauric  acid,  450. 

Chlorocuprous  acid,  432. 

Chloroform  from  methane,  287. 

Chloroplatinates  from  amines,  423. 

Chloroplatinic  acid,  preparation,  565  ; 
from  platinum  black,  566 ;  use  in 
photography,  445. 

Chloroplatinous  acid,  565. 

Chloroplumbic  acid,  518. 

Chloroplumbous  acid,  518. 

Chlorosulfonic  acid,  preparation, 
properties,  189. 

Chlorous  acid,  127 ;    structure,  130. 

Cholera  from  impure  water  supply,  83. 

Chrome  alum,  preparation,  527. 

Chrome  green,  525. 

Chrome  iron  ore,  524 ;  decomposition 
of,  527. 

Chrome  tanning,  527. 

Chrome  yellow,  527 ;  constituent  of 
chrome  green,  525. 


Chromic  anhydride,  preparation,  use 
to  oxidize  carbon,  528. 

Chromic  chloride,  hydrates,  525; 
isomeric,  526  ;  preparation,  proper- 
ties, 525  ;  theory  of,  526. 

Chromic  hydroxide,  formation,  com- 
position, properties,  525 ;  chro- 
mites,  525. 

Chromic  oxide,  preparation,  use  as 
pigment,  525. 

Chromite,  524. 

Chromites,  525. 

Chromium,  alloy  with  cobalt,  557; 
occurrence,  metallurgy,  properties, 
uses,  524 ;  preparation  by  thermite 
process,  524,  498. 

Chromium  trioxide,  preparation,  use 
to  oxidize  carbon,  528. 

Chromous  chloride,  preparation, 
properties,  525. 

Chromyl  chloride,  preparation,  prop- 
erties, structure,  hydrolysis  to  di- 
chromic acid,  528. 

Chugaev,  separation  of  nickel  and 
cobalt,  561. 

Cinnabar,  485,  489. 

Citral,  use  in  making  ionpne,  328. 

Citrate-soluble  phosphoric  acid,  461. 

Citric  acid,  structure,  source,  330 ; 
salts,  331. 

Clark  cell,  electromotive  force,  437. 

Clarke,  F.  W.  Composition  of  the 
crust  of  the  earth,  11. 

Classification,  of  metals,  370,  371 ; 
of  the  elements,  132. 

Clays,  formation  of,  494;  manufac- 
ture of  aluminium  oxide,  hydro- 
chloric acid  and  sodium  carbonate 
from,  with  salt,  496. 

Cleveite,  discovery  of  helium  in,  237 ; 
uranium  in,  531. 

Clinker,  cement,  454. 

Clouds,  conditions  of  formation,  232. 

Clover,  fixation  of  nitrogen  by,  199. 

Coal,  energy  of  from  sunlight,  230 ; 
formation,  varieties,  composition, 
280. 

Coal  gas,  percentage  composition, 
299  ;  tar,  character,  294 ;  tar  dips, 
phenol  in,  326. 

Coals,  coking,  noncoking  and  cannel, 
281. 

Cobalt  ammines,  559 ;  glass,  smalt, 
558 ;  nitrate,  use  in  blowpipe 
analysis,  558. 

Cobalt,  occurrence,  properties,  alloy 
with  chromium  oxides,  557 ;  sepa- 
ration from  nickel  by  dimethyl- 
glyoxime,  560  ;  sulfide,  formation, 
slow  solubility,  558. 

Cobaltous  chloride,  properties,  sym- 


INDEX 


577 


pathetic  ink,  557 ;  hydroxide,  prep- 
aration, properties,  557. 

Cocaine,  343  ;  similar  synthetic  alka- 
loids, 343. 

Coefficients  of  expansion  of  air,  O2, 
N2,  NO,  H2,  A,  He,  CO,  CO2,  SO2, 38. 

Coining  value  of  gold,  447. 

Coins,  gold,  448 ;  nickel  five  cent, 
560;  silver,  442. 

Coke,  composition,  280 ;  manufac- 
ture, uses,  278. 

Coking  coals,  281. 

Colemanite,  boric  acid  from,  365. 

Collodion,  338. 

Collection  and  storage  of  gases,  22. 

Colloidal  arsenic  trisulfide,  261 ; 
silicic  acid,  353  ;  solutions,  262. 

Colloids,  contrasted  with  crystalloids, 
357  ;  precipitated  by  bivalent  ions, 
263 ;  properties,  262. 

Columbite,  523. 

Columbium  (niobium),  discovery,  oc- 
currence, properties,  compounds, 
523. 

Combining  volumes,  law  of,  89 ; 
weights,  law  of,  13. 

Combustion,  24  ;  heat  of,  25. 

Complex  cyanides,  319. 

Complex    ions,    formation   of,    378 ; 

'    evidence  for  existence  of,  379. 

Composition  of  air  demonstrated  by 
Lavoisier,  19 ;  of  the  crust  of  the 
earth,  11 ;  of  pure  substances 
expressed  in  multiples  of  atomic 
weights,  17. 

Compounds,  definition,  9 ;  general 
methods  of  preparing,  372-379. 

Compressed  gases,  cooling  on  ex- 
pansion, 233. 

Comstock,  dependence  of  mass  on 
velocity,  5. 

Concentration  and  speed  of  reaction, 
149 ;  effect  of  on  chemical  reac- 
tions, 24. 

Congress  water,  309. 

Coniine,  342. 

Conservation  of  energy,  6 ;  of  matter, 
6. 

Constant  proportion,  law  of,  12. 
Contact     mass "     for     catalyzing 
formation  of  sulfur  trioxide,   176. 

Converter,  Bessemer,  547. 

Cooking  of  starchy  foods,  336. 

Copper,  addition  to  steel,  431 ;  action 
of  nitric  acid  on,  213  ;  acetylide, 
see  copper  carbide. 

Copper,  alloys  of,  431 ;  annual  pro- 
duction, value,  430 ;  properties, 
effect  of  impurities  on  conduc- 
tance of,  430;  uses,  431 ;  arsenite 
and  acetate  (Paris  green),  259. 


Copper,  basic  carbonate  of,  428; 
formation,  431. 

Copper  carbide,  preparation,  292 ; 
chloride,  catalyzer  for  Deacon 
process,  103 ;  detected  by  sodium 
metaphosphate,  253 ;  electro- 
lytic refining,  429. 

Copper,  exercises,  450  ;  ferrocyanide, 
use  in  semipermeable  membranes, 
358  ;  hydroxide,  precipitation,  de- 
composition, 431 ;  in  five  cent 
piece,  560 ;  metallurgy  of,  428. 

Copper,  occurrence,  metallurgy,  428  ; 
electrolytic  refining,  429. 

Copper  oxide,  from  copper  hydroxide, 
431 ;  from  wire,  nitrate,  use,  432  ; 
in  "  oxone,"  21 ;  use  in  deter- 
mination of  the  composition  of 
water,  69. 

Copper,  precipitation  by  sodium 
thiosulfate,  409 ;  precipitation  by 
iron,  435 ;  prevention  of  corro- 
sion of  iron  by,  550;  pyrites,  161, 
428 ;  separation  from  cadium  as 
cuprocyanide,  435. 

Copper,  sulfate,  hydrates,  uses,  433  ; 
reaction  with  sulfuric  acid,  173 ; 
titration  of  with  potassium  iodide 
and  sodium  thiosulfate,  433. 

Copperas,  554. 

Corn  sirup,  334. 

Cornwall,  tin  from,  508. 

Corpuscle,  same  as  electron,   181. 

Correction  of  volume  of  gas  for 
temperature,  39 ;  zero  and  stem, 
for  thermometers,  486. 

Corrections  for  readings  of  barom- 
eter for  altitude,  37 ;  for  read- 
ings of  barometer  for  glass  and 
brass  scales,  36 ;  for  readings  of 
barometer  for  latitude,  37. 

Corrosion  of  iron,  prevention  by 
copper,  550. 

Corrosive  sublimate,  preparation, 
489 ;  properties,  uses,  antidote 
for,  490. 

Cort,  invention  of  puddling  process, 
544. 

Cotton  goods,  fire-proofing  of,  513. 

Cowles  Brothers,  electric  furnace, 
495 ;  use  of  electric  furnace  for 
aluminium,  391. 

Cowles,  process  for  decomposition 
of  clay,  496. 

Cowper-Cowles,  Sherard,  sherard- 
ized  iron,  482. 

Cream  of  tartar,  tartar  emetic  from, 
266 ;  use  in  jellies,  in  baking  pow- 
ders, 330. 

Cretinism,  connected  with  defi- 
ciency of  iodine,  144. 


578 


INDEX 


Critical  temperature,  defined,  233; 
discovery  of,  232 ;  of  carbon 
dioxide,  233 ;  relation  to  lique- 
faction of  air,  233. 

Crookes,  discharge  of  electricity 
through  rarefied  gases,  471 ;  dis- 
covery of  thallium,  507. 

Crown  glass,  467. 

Cryolite,  153,  495 ;  use  for  manu- 
facture of  aluminium,  495. 

Crystallization,  8 ;    water  of,  82. 

Crystallographic  systems,  193. 

Crystalloids,  357. 

Crystals,  definition  of,  192. 

Cubic  centimeter,  true  volume  of,  31. 

Cupellation,  440. 

Cupric  chloride,  preparation,  ioniza- 
tion,  432;  nitrate,  hydrates,  de- 
composition of,  434 ;  oxide,  prep- 
aration, use,  432 ;  sulfide,  forma- 
tion, properties,  433. 

Cuprous  chloride,  preparation,  prop- 
erties, 432;  use  to  absorb  car- 
bon monoxide,  433 ;  cyanide,  for- 
mation, complex  salt  with  potas- 
sium cyanide,  434;  iodide,  forma- 
tion in  titrating  copper,  433; 
oxide,  formation,  in  testing  for 
glucose,  hydrazine,  etc.,  proper- 
ties, 432;  sulfide  in  matte,  429; 
sulfide,  occurrence,  428 ;  forma- 
tion, 433. 

Cyanides,  complex,  319;  formation, 
preparation,  319;  from  ammoni- 
acal  gas  liquors,  319. 

Cyanide  process,  for  gold,  446; 
for  silver,  441. 

Cyclopentene,  structure,  285. 

Cyclopropane,  structure,  284. 

Dalton,  atomic  theory,  14 ;  formula 
for  water,  91 ;  law  of  multiple 
proportion,  88;  law  of  partial 
pressures,  41,  77 ;  view  of  mole- 
cules of  the  elements,  93. 

Damascus  blade,  tungsten  in,  530. 

Davy,  discovery  of  metallic  potassium, 
415 ;  discovery  of  metallic  sodium 
and  potassium,  399;  injured  by 
nitrogen  trichloride,  224;  safety 
lamp,  287 ;  study  of  explosion  of 
fire  damp,  287. 

Deacon  process  for  chlorine,  102; 
equilibrium  of,  109. 

Debye,  quantum  theory,  398. 

Degree  of  ionization,  measurement  of , 
380 ;  table,  acids,  383  ;  bases,  383  • 
salts,  384. 

Degrees  of  freedom,  77. 

Dehydration  by  sulfuric  acid,  182. 

Dekahydronaphthalene,  285. 


Deliquescence,  82. 

Density,  criterion  of  pure  substance, 
12 ;  of  gases,  table,  95. 

Dental  cement,  composition,  482. 

Derivatives  of  ammonia,  205. 

Determination  of  weight  of  a  liter  of 
gas,  40. 

Detonating  caps,  491. 

Developing  in  photography,  445. 

Dewar  flasks  for  liquid  air,  235. 

Dextrin,  manufacture  from  starch, 
uses,  336. 

Dextrose,  see  Glucose. 

Dialysis,  357. 

Dialyzed  iron,  555. 

Diamond,  heat  of  combustion,  276 ; 
artificial,  275,  natural,  275;  uses, 
properties,  275. 

Diastase,  344. 

Dibasic  acids,  defined,  183. 

Dibromoethane,  291. 

Dicalcium  phosphate,  249. 

Dichloroethane,  291. 

Dichromic  acid,  from  hydrolysis  of 
chromyl  chloride,  528. 

Diet,  salt  essential  in,  406. 

Dietary,  average  American,  347. 

Diffusion  of  gases,  56 ;   law  of,  59. 

Digestion,  colloidal  solutions,  263. 

Dimethylglyoxime,  precipitant  for. 
nickel,  560. 

Diphosphorus  pentasulfide,  254. 

Disilicates,  356. 

Disilicic  acid,  355. 

Disinfectant,  formaldehyde,  327  ; 
sulfur  dioxide,  174. 

Disintegration  of  atoms,  possible  by 
radium  emanation,  475. 

Disodium  phosphate,  249 ;  hydroly- 
sis, alkaline  reaction  of,  251 ; 
phenolphthalein  as  indicator  for, 
251 ;  uses,  arsenic  as  impurity  in, 
410. 

Displacement  of  equilibrium,  152. 

Dissociation,  definition,  59  ;  pressures 
of  calcium  carbonate,  453 ;  pres- 
sure of  silver  oxide,  443 ;  of  am- 
monium hydroxide  to  ammonia 
and  water,  204 ;  of  calcium  car- 
bonate, phase  rule,  453  ;  of  sulfuric 
acid,  180 ;  of  water,  59. 

Distillation,  8. 

Dithionic  acid,  188. 

Divariant,  definition,  77. 

Dixon,  explosion  waves,  301. 

Dolomite,  478. 

Double  decomposition,  defined,  81. 

Double  refraction  of  crystals,  196. 

Dry  batteries,  use  of  zinc  in,  482. 

Dry  plates,  photographic,  444. 

Drying  of  gases,  54. 


INDEX 


579 


Dulong  and  Petit,  law  of,  396; 
relation  to  Ayogadro's  law,  397. 

Dumas,  determination  of  the  compo- 
sition of  water,  69. 

Durax  glass,  467. 

Dust,  explosion  of  with  air,  289. 

Dutch  process  for  white  lead,  520. 

Dydimium,  separation  into  praseo- 
dymium and  neodymium,  504. 

Dyeing,  lakes  for,  501. 

Dyes,  340 ;  use  of  sodium  nitrite  in 
manufacture  of,  410  ;  substantive, 
adjective,  342. 

Dysprosium,  compounds,  505. 

Earth,  composition  of  crust  of,  11 ; 
mean  density  of,  540. 

Earthenware,  501,  glazing,  502. 

Effervescent  waters,  309. 

Efflorescence,  82. 

Egypt,  early  manufacture  of  iron  in, 
390  ;  sodium  carbonate  from,  450. 

Einstein,  quantum  theory,  398. 

Eka-aluminium,  same  as  gallium,  506. 

Ekaboron,  same  as  scandium,  136, 503. 

Ekeberg,  discovery  of  tantalite,  523. 

Electric  furnace,  carborundum  in, 
349 ;  use  for  aluminium,  391. 

Electrical  batteries,  use  of  zinc  in, 
481,  theory  of,  435 ;  horse  power, 
34  ;  unit  charge,  438  ;  units,  33. 

Electrochemical  theory,  influence  on 
formulas  of  minerals,  356. 

Electrodes,  carbon,  279. 

Electrolysis,  migration  of  ions  in,  113  ; 
of  dilute  sulfuric  acid,  47,  9 ; 
sodium  hydroxide  by,  402. 

Electrolyte,  definition,  48. 

Electrolytic  methods  in  metallurgy, 
391. 

Electromotive  force,  relation  to  solu- 
tion pressure,  435. 

Electromotive  series,  435,  table,  436. 

Electron,  relation  to  Faraday's  law, 
438. 

Electron  theory,  181 ;  as  explanation 
of  ionization,  182 ;  relation  to 
ionization,  206;  relation  to  prop- 
erties of  metals  and  non-metals, 
370. 

Electronegative  elements,  denned, 
437. 

Electroplating,  copper,  433. 

Electropositive  elements,  defined, 
437. 

Electrotyping,  433. 

Elements,  absolute  potential  of,  436 ; 
atoms  of  probably  complex  aggre- 
gates, 138  ;  classification  of,  132  ; 
definition,  9  ;  life  of,  474  ;  melting 
points  of,  absolute,  135;  melting 


points  of,  372,  table,  373  ;  metallic 
in  periodic  system,  136 ;  missing 
in  Group  VII,  possible  reason,  533  ; 
molecules  of,  93. 

Elements,  non-metallic  in  periodic 
system,  136 ;  radioactive,  series  of, 
475 ;  specific  heat  of,  397  ;  symbols 
of,  11;  table,  of  familiar,  371; 
table  of  groups  of,  371 ;  table 
of  non-metallic,  348. 

Emery,  494;  artificial,  500. 

Endothermic  compounds,  defined, 
225 ;  explosive  decomposition,  225. 

Endothermic  reactions,  215. 

Energy,  conservation  of,  6 ;  defini- 
tion, 6;  muscular,  in  respiration 
calorimeter,  315;  of  coal  from 
sunlight,  230. 

Energy  required  to  decompose  a 
gram  equivalent,  438  ;  units  of,  32 ; 
varieties  of,  6. 

England,  destruction  of  forests  for 
iron  manufacture,  540. 

English  laws  for  glazes,  502. 

Enzymes,  344. 

Epsom  salts,  478,  480. 

Equations,  writing  of,  49 ;  writing  of, 
for  reactions  between  acids  and 
bases,  156. 

Equilibrium  between  gaseous  and 
solid  phases,  443,  silver,  silver 
oxide  and  oxygen,  443 ;  between 
water  and  water  vapor,  76;  dis- 
placement of,  152 ;  effect  on,  of 
removing  one  of  the  reacting  sub- 
stances, 152  ;  for  combination  of 
nitrogen  and  hydrogen,  201 ;  for 
formation  of  nitric  oxide,  216; 
hydrogen,  iodine  and  hydriodic 
acid,  146. 

Equilibrium  in  chemical  reactions, 
108 ;  in  gas  flame,  300 ;  illustra- 
tion of  that  between  water  and 
water  vapor,  76 ;  in  ionization  of 
orthophosphoric  acid,  250 ;  in 
neutralization,  385. 

Equilibrium  of  carbon  monoxide, 
carbon  dioxide,  hydrogen  and  water 
vapor,  300 ;  of  Deacon  process, 
109 ;  of  reaction  between  hydro- 
chloric acid  and  oxygen,  108. 

Equivalents,  relation  to  Faraday's 
law,  438. 

Erbium,  compounds,  505. 

Erg,  defined,  32. 

Ethane,  structure,  284;  substitu- 
tion products  of,  291. 

Ethene,  290. 

Ether,  ethyl,  290. 

Ethyl  alcohol,  ether  and  ethylene 
from,  290;  ethyl  chloride  from, 


580 


INDEX 


245 ;  manufacture,  properties,  uses, 
325;  absolute,  denatured,  325; 
structure,  323. 

Ethyl  borate,  367;  chloride,  formed 
from  ethyl  alcohol,  245;  ether, 
formation,  290;  iodide,  relation 
to  structure  of  ethyl  alcohol,  323. 

Ethylene,  addition  compounds,  291 ; 
bromide,  291 ;  chloride,  formation 
from  ethylene,  291 ;  in  illuminat- 
ing gas,  295 ;  preparation,  prop- 
erties, uses,  formation  and  de- 
composition, 290  ;  ratio  of  specific 
heats  of,  237 ;  structure,  292. 

Eudiometer,  description,  67. 

Europium,  compounds,  505. 

Eutectic  point,  defined,  488. 

Evaporators,  triple  and  multiple- 
effect,  405 ;  Yaryan,  405. 

Exercises,  aluminium,  507 ;  atmos- 
phere, 239;  Avogadro's  law,  99; 
calcium,  barium,  476 ;  carbon 
dioxide,  cyanides,  322 ;  chlorine, 
116;  copper,  silver,  gold,  450. 

Exercises,  Group  V,  272  ;  hydrocar- 
bons, 305 ;  laws  of  gases,  43  ;  mag- 
nesium and  mercury,  493 ;  nitro- 
gen, 225  ;  phosphorus,  255 ;  silicon, 
boron,  368;  sulfur,  196;  writing 
equations,  156. 

Exothermic  reaction,  definition  of, 
215. 

Explosions,  definition,  62 ;  of  en- 
dothermic  compounds,  225 ;  of 
methane  and  air,  287;  of  dust  and 
air,  289  ;  waves,  301. 

Factories,  humidity  in,  232. 

Families  of  elements,  table,  371. 

Faraday,  injured  by  nitrogen  tri- 
chloride, 224;  law,  338;  liquefac- 
tion of  chlorine,  108. 

Fast  colors,  340. 

Fats,  composition,  331 ;  use  in 
making  soaps,  332  ;  reduction  with 
the  aid  of  colloidal  palladium,  564  ; 
soft  soap  from,  414. 

Fatty  acids,  calcium  salts  of,  332. 

Fehling's  solution,  formula  for,  use, 
to  detect  glucose,  335. 

Feldspars,  348 ;    use  as  glaze,  502. 

Ferric  acetate,  mordant,  342. 

Ferric  chloride,  hydrate,  anhydrous, 
554 ;  molecular  weight,  555 ; 
hydrolysis,  precipitation  of  ferric 
hydroxide  from,  553;  theory  of 
hydrolysis  of,  386. 

Ferric  ferrocyanide,  320 ;  decomposi- 
tion with  sodium  hydroxide,  321 ; 
hydroxide,  precipitation,  555;  by 
barium  carbonate,  553  ;  hydroxide, 


to  absorb  hydrogen  sulfide,  295; 
oxide,  preparation,  manufacture, 
use,  555. 

Ferric  sulfate,  fuming  sulfuric  acid 
from,  556 ;  preparation,  alums  from, 
556 ;  reduction  by  hydrogen  sul- 
fide, 171 ;  sulfide,  formation,  prop- 
erties, 556 ;  thiocyanate,  test  for 
iron,  322  ;  thiocyanate,  use  as  test 
for  iron,  556. 

Ferrite,  defined,  a-,  /3-  and  7-,  rela- 
tion to  tempering,  546 ;  table  of 
properties,  546. 

Ferromanganese,  use  in  chilled  cast 
iron,  543. 

Ferrous  bicarbonate,  formation,  in 
mineral  waters,  ores  from,  554 ; 
carbonate,  ore,  540 ;  carbonate, 
properties,  554. 

Ferrous  chloride,  absorption  of  nitric 
oxide  by  solutions  of,  554  ;  prepara- 
tion, properties,  553  ;  preparation, 
structure,  552. 

Ferrous  chromite,  524 ;  preparation 
of  potassium  chromate  from,  527 ; 
ferricyanide,  321 ;  hydroxide,  for- 
mation, properties,  553  ;  hydroxide, 
use  to  reduce  indigo,  341 ;  manga- 
nese tungstate,  530 ;  metatantalate, 
523  ;  metacolumbate,  523  ;  oxide, 
553  ;  silicate,  formed  in  metallurgy 
of  copper,  429. 

Ferrous  sulfate,  preparation,  proper- 
ties, 554 ;  sulfide,  oxidation  to 
ferric  hydroxide  in  gas  purifiers, 
296 ;  sulfide,  preparation,  prop- 
erties, use,  556. 

Ferro vanadium,  522. 

Ferrum,  11. 

Fertilizer,  calcium  cyanamide  as,  463 ; 
slag  from  basic  steel  process  as,  548  ; 
various  phosphates  in,  461. 

Filters,  charcoal,  inefficient,  83. 

Fire  damp,  286,  287. 

Fireproofing  fabrics,  with  sodium 
tungstate,  530 ;  of  cotton  goods, 
513  ;  with  water  glass,  353. 

Fixing  in  photography,  445. 

Flame  colors  of  calcium,  strontium 
and  barium,  471 ;  gas  and  candle, 
structure,  299. 

Flame,  oxidizing,  blowpipe  and  Bun- 
sen  burner,  304;  reducing,  blow- 
pipe and  Bunsen  burner,  304. 

Flames,  cause  of  luminosity,  299  ;  re- 
versed, 304  ;  temperature  of,  302  ; 
source  of  carbon  in,  299. 

Flint,  348. 

Flint,  atomic  weight  of  tellurium, 
190. 

Flint  glass,  467. 


INDEX 


581 


Flour,  explosive  when  mixed  with  air, 
289. 

Flowers  of  sulfur,  160. 

Fluoboric  acid,  367. 

Fluorine,  occurrence,  153 ;  prepara- 
tion, 153 ;  properties,  154. 

Fluorite,  153,  452. 

Fluosilicic  acid,  preparation,  proper- 
ties, 350. 

Foods,  animal,  vegetable,  inorganic 
constituents,  347  ;  average  Ameri- 
can dietary,  347. 

"  Fool's  gold,"  556. 

Forge,  reduction  of  iron  in,  540. 

Formaldehyde,  preparation,  use  as 
disinfectant,  327 ;  use  in  food  for- 
bidden, 328 ;  probable  formation 
by  reduction  of  carbon  dioxide  by 
plants,  312. 

Formalin,  see  formaldehyde. 

Formic  acid,  from  mercuric  fulminate, 
491 ;  preparation  from  carbon 
monoxide,  uses,  329. 

Formula  of  mineral,  calculation,  357. 

Formular  solutions,  definition  of, 
183. 

Formulas,  meaning  of,  16;  struc- 
tural, basis  for,  323. 

Franklinite,  481. 

Frasch,  process  for  getting  sulfur,  161. 

Fraunhofer  lines,  discovery,  424. 

Free  metals,  occurrence,  439. 

Freedom,  degrees  of,  77. 

Freezing  points  of  solutions  and  os- 
motic pressure,  360 ;  law  for  de- 
pression of,  112. 

French  Revolution,  relation  to  alkali 
industry,  400  ;  sources  of  saltpeter 
during,  418. 

Friedel,  discovery  of  carbon  in  meteo- 
rite from  Canon  Diablo,  274. 

Fructose,  formation  from  cane  sugar, 
properties,  335;  from  cane  sugar, 
334. 

Fruit,  sulfuring  of,  164. 

Furnace,  electric  for  aluminium,  495  ; 
electric  for  calcium  carbide,  462 ; 
electric  for  carbon  disulfide,  317; 
electric  for  carborundum,  349 ; 
electric  for  phosphorus,  241 ;  re- 
generative, for  open  hearth  steel, 
549,  550. 

Fusible  alloys,  269. 

Gadolinium,  compounds,  505. 

Galena,  161 ;  lead  ore,  513  ;  silver  in, 
439. 

Gallium,  compounds,  same  as  eka- 
aluminium,  506. 

Galvanic  cells,  calculation  of  electro- 
motive force,  437 ;  theory,  437. 


Galvanized  iron,  manufacture,  theory 
of  conduct,  482. 

Garnet,  349 ;  calculation  of  formula 
of,  356  ;  orthosilicate,  355. 

Gas,  blast  furnace ;  use  for  heating 
blast,  for  gas  engines,  542 ;  carbon, 
manufacture,  uses,  279. 

Gas,  correction  of  volume  for  changes 
of  pressure,  38 ;  for  temperature,  39. 

Gas,  determination  of  weight  of  a 
liter  of,  40 ;  effect  of  pressure  on, 
34 ;  effect  of  temperature  on,  38 ; 
effect  of  water  vapor  on  volume 
and  pressure  of,  76. 

Gas,  illuminating,  295 ;  "  ideal,"  94  ; 
iron  pentacarbonyl  in,  561 ;  laws, 
graphical  representation,  42,  43 ; 
lighters,  cerium-iron,  364 ;  liquors, 
ammoniacal,  202. 

Gas,  oil,  Pintsch,  water,  296 ;  pro- 
ducer, 297 ;  producer,  use  in  re- 
generative furnace  for  steel,  550. 

Gases,  collection  and  storage  of,  22 ; 
diffusion  of,  56 ;  drying  of,  54 ; 
Henry's  law  for  solution  of,  165 ; 
spectra  of,  427  ;  kinetic  theory  of, 
58 ;  laws  of,  34,  38  ;  laws  of,  exer- 
cises, 43 ;  monatomic,  ratio  of 
specific  heats  for,  236;  table  of 
densities,  95. 

Gasoline,  289. 

Gasometer,  description  of,  22. 

Gastric  juice,  hydrochloric  acid  in, 
406. 

Gayley,  dry  blast  for  iron  manufac- 
ture, 542. 

Gay  Lussac's  law  of  combining 
volumes,  89  ;  tower,  use  in  manu- 
facture of  sulfuric  acid,  179. 

Geology,  relation  of  radiochemistry 
to,  475. 

Germanium,  discovery,  compounds, 
361. 

German  silver,  560. 

Germicides,  sulfites  as,  174. 

Gibbs,  Willard,  phase  rule,  107. 

Glass,  etching  of,  154  ;  manufacture, 
properties,  466 ;  crown,  flint, 
strass,  paste,  Bohemian,  Jena,  Re- 
sistanz,  Non-sol,  hard,  borosilicate, 
Durax,  467  ;  opaque,  stannic  oxide 
in,  511 ;  soluble  or  water,  353. 

Glazes,  for  earthenware  and  porce- 
lain, 502. 

Glover  tower,  use  in  sulfuric  acid 
manufacture,  179. 

Glucinum,  name  for  beryllium,  451. 

Glucose,  from  hydrolysis  of  cane 
sugar  and  of  starch,  properties, 
334 ;  formation  in  diabetes,  defec- 
tion, 335. 


582 


INDEX 


Gluten  from  cereals,  8,  335. 

Glycerol,  source,  use,  326;  by-prod- 
uct in  manufacture  of  soap,  322. 

Glyceryl  nitrate,  326. 

Gmelin,  manufacture  of  ultramarine, 
502. 

Goiter,  connected  with  deficiency  of 
iodine,  144. 

Gold,  chlorides  of,  450 ;  dioxide,  448 ; 
exercises,  450 ;  monochloride,  450  ; 
dichloride,  450  ;  trichloride,  45 ; 
monoxide,  448. 

Gold,  occurrence,  445 ;  in  sea  water, 
washing  for,  hydraulic  mining, 
cyanide  process,  446  ;  annual  pro- 
duction, coining  value,  447  ;  oxides 
of,  hydroxide,  448;  properties, 
alloys,  use  of  term  carat,  448. 

Gold,  recovery  from  copper,  430 ; 
trichloride,  use  in  photography, 
445;  trioxide,  448. 

Goldschmidt's  thermite  process,  497. 

Goldthwaite,  jelly  making,  337. 

Goodwin,  preparation  of  metallic 
calcium,  452. 

Graebe,  discovery  of  structure  of 
alizarin,  341. 

Graham,  crystalloids,  colloids,  357. 

Graham,  study  of  diffusion  of  gases, 
57 ;  of  liquids,  357. 

Gram,  definition,  31. 

Gram  atom,  definition,  25. 

Gram  equivalent,  definition  of,  184. 

Gram  molecular  volume,  definition, 
94. 

Gram  molecule,  definition,  25. 

Granites,  decomposition  to  shales, 
clays  and  soils,  494. 

Graphical  representation  of  gas  laws, 
42,  43. 

Graphite,  abnormal  specific  heat, 
397 ;  anodes  for  alkali  manufac- 
ture, 402  ;  heat  of  combustion,  276 ; 
in  cast  iron,  543,  544 ;  in  colloidal 
solution  as  lubricant,  277  ;  sources, 
formation,  manufacture,  proper- 
ties, 276 ;  uses,  276. 

Gravitation,  law  of,  1. 

Gravity  cell,  description  of,  electro- 
motive force  of,  437. 

Greenland,  cryolite  from,  495. 

Green  vitriol,  554. 

Grotto  del  Cano,  carbon  dioxide  in, 
308. 

Groups  of  elements,  table  of,  371. 

Group  zero,  236. 

Group  I,  alternate  metals  of,  general 
properties,  428. 

Group  II,  alternate  metals  of,  general 
properties,  478 ;  general  properties, 
451 ;  solubility  of  sulfides  of,  491. 


Group  III,  metals  of,  494. 

Group  IV,  metallic  elements  of,  361, 
362. 

Group  V,  alternate  elements,  522; 
exercises,  272  ;  general  properties, 
256 ;  table  of  oxides,  chlorides, 
acids,  sulfur  acids  and  hydrides, 
271. 

Group  VI,  alternate  elements,  524; 
metallic  elements  of,  192 ;  proper- 
ties of  elements  of,  191. 

Group  VII,  metallic  elements  of,  156, 
533 ;  missing  elements,  possible 
reason,  533. 

Group  VIII,  general  statement,  539 ; 
metals  of,  general  properties,  table, 
562. 

Guignet's  green,  525. 

Gun  cotton  in  smokeless  powder,  419  ; 
manufacture,  properties,  uses,  338. 

Gun  metal,  509. 

Gunpowder,  composition,  418 ;  de- 
composition, manufacture,  theory 
of  burning  and  explosion,  419. 

Gypsum,  161,  452;  plaster  of  Paris 
from,  457 ;  phase  rule,  458 ;  con- 
ditions for  formation  of,  458,  460 ; 
vapor  pressure  of,  table,  459. 

Haber,  dissociation  of  hydrobromic 
acid,  146 ;  synthesis  of  ammonia, 
201. 

Haddam,  columbite  from,  523. 

Hall,  manufacture  of  aluminium,  495  ; 
process,  496. 

Halogen  acids,  table  of,  139 ;  family, 
139  ;  meaning  of  name,  139. 

Halogens,  compounds  with  hydrogen 
and  oxygen,  139 ;  general  prop- 
erties of,  138;  table  of,  139. 

Hamburg,  cholera  in,  from  water, 
83. 

Hampson,  liquid  air  machine,  234. 

Harcourt,  atomic  weight  of  tellu- 
rium, 190. 

Hard  glass,  467. 

Hard  waters,  310 ;  permanent  hard- 
ness, temporary  hardness,  soften- 
ing by  boiling,  463  ;  by  milk  of  lime, 
464 ;  by  sodium  salts,  464. 

Hatchett,  discovery  of  columbium, 
523. 

Haynes,  cobalt  chromium  alloy,  557. 

Heat  of  combination  of  chlorine 
with  H2,  Na,  Zn,  Cu,  P,  108; 
of  oxygen  with  Ek,  Na,  Zn,  Cu,  P, 
108. 

Heat  of  combustion,  25 ;  of  bitumin- 
ous coal,  calculation  of,  44 ;  of 
C,  S,  P,  Fe,  Hg,  27 ;  of  hydrogen, 
65. 


INDEX 


583 


Heat,  mechanical  equivalent  of,  6 ; 
units  of,  33. 

Helium,  coefficient  of  expansion,  38 ; 
discovery  in  sun  and  in  cleveite, 
237;  distribution,  238;  from 
radium,  9. 

Hematite,  540. 

Hemihedral  forms  of  crystals,  193. 

Henry's  law,  applied  to  carbon  di- 
oxide, 309 ;  for  solution  of  gases, 
165. 

Heptane,  283. 

Heptene,  283. 

Heptine,  283. 

Hexaaquochromic  chloride,  526. 

Hexagonal  System  (crystallography), 
194. 

Hexane,  283. 

Hexanitrocellulose,  338. 

Hexathionic  acid,  188. 

Hexene,  283. 

Hexine,  283. 

Hillebrand,  gases  from  uraninite, 
237 ;  use  of  sodium  pyrosulfate 
for  solution  of  alumina,  408. 

Holmberg,  holmium,  505. 

Holmium  source,  505. 

Homologue,  definition,  289. 

Honey,  invert  sugar  in,  334. 

Hornblende,    metasilicate,    355. 

Horn  silver,  439. 

Hulett,  purification  of  mercury, 
485. 

Humidity  of  rooms,  232. 

Hunyadi  water,  magnesium  sulfate 
in,  480. 

Hydrargyrum,  11. 

Hydrates,  defined,  81 ;  of  nitric 
acid,  211 ;  of  sulfuric  acid,  181 ; 
vapor  pressure  of,  82. 

Hydration,  water  of,  82. 

Hydrazine  hydrochloride,  222 ;  prep- 
aration, 222;  structure,  220; 
trinitride,  structure,  221. 

Hydriodic  acid,  constant  boiling 
solution,  146 ;  equilibrium  with 
hydrogen  and  iodine,  146  ;  forma- 
tion from  potassium  iodide,  145; 
heat  of  formation  of,  153 ;  prep- 
aration from  iodine,  phosphorus 
and  water,  145 ;  rate  of  formation 
and  decomposition,  150 ;  reduc- 
tion of  sulfuric  acid  by,  145. 

Hydrobromic  acid,  constant  boiling 
solution  of,  143 ;  dissociation  of, 
146 ;  preparation  from  hydrogen 
and  bromine,  143 ;  preparation 
from  potassium  bromide,  142 ; 
reduction  of  sulfuric  acid  by,  142 ; 
preparation  from  bromine,  phos- 
phorus and  water,  143. 


Hydrocarbonate  ion,  r61e  in  decom- 
position of  carbonates,  375. 

Hydrocarbons,  exercises,  305 ;  table 
of,  283. 

Hydrochloric  acid,  and  oxygen, 
equilibrium  of  reaction  between, 
108 ;  by-product  in  making  alumin- 
ium oxide  from  clay,  496 ;  diffi- 
culty with  in  Leblanc  soda  process, 
411. 

Hydrochloric  acid,  constant  boiling 
solution  of,  120. 

Hydrochloric  acid,  determination  of 
composition  by  volume,  120 ;  de- 
termination of  composition  of, 
130 ;  deviation  from  Boyle's  law, 
35 ;  electrolysis  to  show  composi- 
tion by  volume,  120 ;  formation 
from  chlorine  and  hydrogen,  118; 
in  gastric  juice,  406 ;  oxidation  of, 
100. 

Hydrochloric  acid,  preparation  from 
salt  and  sulfuric  acid,  118;  prop- 
erties, 119;  reactions  with  hy- 
droxides and  oxides,  121 ;  re- 
actions with  metals,  120 ;  reac- 
tions with  oxidizing  agents,  122; 
reaction  with  potassium  perman- 
ganate, 159. 

Hydrocyanic  acid,  preparation,  prop- 
erties, uses,  319. 

Hydrofluoric  acid,  constant  boiling 
solution  of,  155 ;  etching  of  glass 
by,  154 ;  molecular  weight,  155 ; 
preparation,  154 ;  properties,  154. 

Hydrogen,  45 ;  apparatus  for  prep- 
aration of,  53 ;  burning  in  chlo- 
rine, 118;  chemical  properties  of, 
59 ;  coefficient  of  expansion,  38. 

Hydrogen,  combination  of  iodine  with 
reversible,  146 ;  combination  with 
oxygen  reversible,  372 ;  deter- 
mination of  atomic  weight,  72 ; 
deviation  from  Boyle's  law,  35 ; 
formed  in  preparing  phosphine,  243. 

Hydrogen,  formerly  unit  for  atomic 
weights,  68 ;  heat  of  combustion 
of,  65 ;  ions  in  liquid  ammonia, 
208 ;  in  the  atmosphere,  45. 

Hydrogen,  occurrence,  45 ;  nascent, 
213 ;  palladium  semipermeable 
membrane  for,  358,  360. 

Hydrogen,  preparation,  47 ;  prepara- 
tion by  "  hy drone,"  52;  prepara- 
tion by  sodium  and  potassium, 
50;  preparation  by  zinc  and 
acids,  52 ;  preparation  from  iron 
and  steam,  49 ;  properties  of,  55 ; 
purification  of,  54. 

Hydrogen  peroxide  as  oxidizing 
agent,  84 ;  as  reducing  agent,  85 ; 


584 


INDEX 


preparation,  83;  properties,  85; 
uses,  85 ;  structure,  86 ;  tests 
for,  86. 

Hydrogen  phosphide,  liquid,  243; 
selenide,  preparation,  properties, 
190;  silicide,  349. 

Hydrogen  sulfide,  decomposition  by 
heat,  164 ;  decomposition  of  solu- 
tion in  air,  166;  formation  from 
elements,  164 ;  ionization  of,  168  ; 
occurrence,  164. 

Hydrogen  sulfide,  Parsons  apparatus 
for  generating,  165  ;  reaction  with 
iodine,  171 ;  reducing  agent,  171 ; 
removal  from  hydrogen,  55 ;  re- 
moval from  illuminating  gas,  295 ; 
solubility  in  water,  166. 

Hydrogen  telluride,  190. 

Hydrolysis  of  chlorides,  115;  of 
chlorides  of  phosphorus,  245 ; 
of  salts,  theory,  385;  of  sulfides, 
171. 

"  Hy drone,"  preparation  of  hydrogen 
by,  52. 

Hydronitric  acid,  preparation,  223 ; 
properties,  223;  structure,  221. 

Hydrosulfate  ion,  role  in  decomposi- 
tion of  salt,  375. 

Hydrosulfites,  see  hyposulfites. 

Hydrosulfuric  acid,  167. 

Hydrosulfurous  acid,  186. 

Hydrotetrachloroantimonic  acid, 

267. 

Hydroxide,  meaning  of  name,  21. 

Hydroxides,  oxides  prepared  from, 
392;  preparation  from  metals, 
392;  from  salts,  393. 

Hydroxylamine,  from  mercuric  ful- 
minate, 491 ;  preparation,  221 ; 
properties,  uses,  222. 

Hydroxylammonium  sulfate,   222. 

Hygrometer,  moist  bulb,  232. 

Hypo-,  prefix,  123. 

Hypobromite,  sodium,  143. 

Hypochlorites,  autoxidation  to  chlo- 
rates, 125. 

Hypochlorites,  preparation,  proper- 
ties, uses,  124. 

Hypochlorous  acid,  preparation, 
properties,  124;  formation  from 
chlorine  and  water,  106 ;  structure, 
130. 

Hypochlorous   anhydride,    126. 

Hyponitrous  acid,  preparation,  prop- 
erties, 221. 

Hypophosphites,  use,  preparation, 
248;  structure,  247. 

Hypophosphoric  acid,  formation, 
salts,  254. 

Hyposulfite,  old  name  for  thiosul- 
fate,  187. 


Hyposulfites,    preparation,    186. 
Hypotheses,  2. 

lanke,  George,  value  of  the  calorie 
at  different  temperatures,  33. 

-ic,  suffix,  meaning,  30 ;  use  for  acids, 
123. 

Ice  machines,  204. 

Ice,  vapor  pressure  of,  75. 

-ide,  suffix,  use,  47. 

Illuminating  gas,  ammonia  from,  201 ; 
manufacture,  composition,  295 ; 
removal  of  hydrogen  sulfide  from, 
295. 

Imide,  definition  of,  206. 

Indestructibility  of  matter,  6. 

India,  diamonds  from,  275 ;  potas- 
sium nitrate  from,  418. 

India  rubber  as  semipermeable  mem- 
brane, 358. 

Indicators,  acidity  or  alkalinity  of 
at  change  of  color,  table,  388 ; 
chemical  nature,  389  ;  choice  of  an, 
389  ;  definition  and  list,  122  ;  for 
determining  free  and  combined 
carbonic  acid,  464 ;  use  of,  387  ; 
for  weak  acids  and  bases,  389. 

Indigo,  source,  synthesis,  use  as  dye, 
341 ;  use  of  sodamide  in  manu- 
facture of,  410. 

Indigo  white,  342. 

Indium,  discovery,  atomic  weight, 
compounds,  506. 

Inductive  reasoning,  13. 

Infusorial  earth,  use  for  dynamite, 
packing  and  scouring,  351. 

Ingle,  separated  Bunsen  flame,  301. 

Ink,  sympathetic,  557. 

Insolubility,  effect  on  a  reaction,  376. 

International  Bureau  of  Weights  and 
Measures,  31 ;  scale  of  tempera- 
Invariant,  definition,  78.  [tures,  32. 

Invert  sugar,  334. 

lodic  acid,  139. 

Iodine  and  starch,  145 ;  com- 
bination with  iodine  reversible, 
146;  in  thyroid  gland,  144; 
monatomic  at  high  temperatures, 
144 ;  occurrence,  144  ;  liberated  by 
nitrous  acid,  145 ;  liberated  in 
titrating  copper,  433 ;  positive  in 
iodine  trinitride,  224. 

Iodine,  properties,  144  ;  reaction  with 
hydrogen  sulfide,  171 ;  removal 
from  hydriodic  acid,  145 ;  sodium 
tetrathionate  formed  by  action  of 
on  thiosulfate,  409;  solutions 
standardized  by  arsenious  oxide, 
260;  tincture  of,  144;  trinitride, 
iodine  positive  in,  223 ;  trinitride, 
preparation,  223. 


INDEX 


585 


Ion,  definition,  48. 

Ionium,  475. 

lonization,  calculation  of  degree  of 
from  freezing  points  of  solutions, 
381. 

lonization,  degree  of,  table  acids, 
383 ;  bases,  383 ;  salts,  384  ;  effect 
of  degree  of,  neutralization,  384 ; 
effect  on  freezing  points  of  solu- 
tions, 112,  381. 

lonization,  evidence  of,  112;  ex- 
plained by  the  electron  theory,  182  ; 
measurement  of  degree  of,  380  ;  of 
acids,  167 ;  of  acids,  relation  to 
hydrolysis  of  sugar,  381 ;  of 
ammonium  hydroxide,  203 ;  of 
compounds  of  cadmium  and  mer- 
cury, 491 ;  of  first  and  second 
hydrogen  atom  of  acids,  168 ;  of 
hydrates  of  chromic  chlorides,  526. 

lonization  of  oxalic  acid,  relation  to 
solubility  of  calcium  oxalate,  465, 
466 ;  of  sulfuric  acid,  181 ;  of 
sulfurous  acid,  174 ;  of  trimethyl 
ammonium  hydroxide  and  tetra- 
methyl  ammonium  hydroxide,  204  ; 
of  water,  171,  383  ;  relation  to  elec- 
tron theory,  206. 

Indium,  properties,  uses,  oxides, 
chlorides,  double  salts,  565 ;  tetra- 
chloride,  565. 

Iron,  burning  in  oxygen,  23. 

Iron  carbide,  543,  relation  to  tem- 
pering of  steel,  546. 

Iron,  dialyzed,  555 ;  discovery  of 
metallurgy  of,  390 ;  disulfide,  556  ; 
heat  of  combustion,  27 ;  in  pro- 
teins, 343 ;  magnetic  oxide  of 
from  burning  iron,  23. 

Iron,  precipitation  of  copper  by,  435  ; 
pentacarbonyl,  561 ;  pyrites,  556. 

Iron,  reasons  for  importance  of,  539 ; 
ores,  history  of  use,  540,  blast 
furnace,  541 ;  cast  iron,  543 ; 
wrought,  544 ;  cementation  and 
cast  steel,  tempering  of  steel,  545 ; 
Bessemer  steel,  547 ;  open  hearth 
steel,  548 ;  analyses,  544,  551 ; 
'alloy  steels,  552  ;  compounds,  552  ; 
tetracarbonyl,  561. 

Isobutane,  structure,  284. 

Isomer,  definition,  323 ;  history  of 
name,  511. 

Isometric  System  (crystallography), 
193. 

Isothermals  of  carbon  dioxide,  307. 

-ite,  use  for  salts,  124. 

Jakowin,  reaction   of   chlorine    with 

water,  106. 
James,  separation  of  thulium,  506. 


Jasper,  348. 

Jelly,  conditions  for  making,  337. 

Jena  glass,  467. 

Johnson,  sherardized  iron,  482. 

Jqule,  defined,  26  ;  relation  to  calorie. 

33. 
Joule-Thomson,  effect  on  expansion 

of  compressed  gases,  233. 

Kalium,  11. 

Kalk-Stickstoff,  463. 

Kamm,  use  of  potassium  silver  co- 
baltinitrite  in  analysis,  559. 

Kaolin,  349;  formation,  494;  ortho- 
silicate,  355. 

Kekule,  formula  for  benzene,  285. 

Kelly,  inventor  of  process  for  steel, 
547. 

Kelvin,  size  of  molecules,  16 ;  theory 
of  the  source  of  oxygen  in  the  air, 
230. 

Kerosene,  manufacture,  flashing 
point,  290. 

Ketones,  327. 

Kieselguhr,  359. 

Kilogram-meter,  defined,  32;  value 
in  ergs,  33. 

Kilowatt,  defined,  34. 

Kimberly,  diamonds  from,  275. 

Kindling  temperature,  24. 

Kinetic  theory  of  gases,  58 ;  rela- 
tion to  osmotic  pressure,  361. 

Kipp  generator,  54. 

Kirchoff,  discovery  of  rubidium  and 
caesium,  424 ;  discovery  of  spec- 
trum analysis,  424. 

Knietsch,  history  of  catalytic  sulfur 
trioxide,  175. 

Konigsberger,  critical  temperature  of 
mercury,  486. 

Kremann,  hydrates  of  nitric  acid,  212. 

Krogh,  estimate  of  coal  burned 
annually,  229. 

Krypton,   discovery,    238. 

Kurnakow,   sodium  amalgams,  488. 

Kiister  and  Kremann,  hydrates  of 
nitric  acid,  212. 

Lacquers,  use  of  nitrocellulose  in,  338. 

Lactic  acid,  formation,  structure,  330. 

Lactose,  from  milk,  use,  334. 

Ladenburg,  density  of  ozone,  97. 

Lake  Superior  region,  copper  from, 
428. 

Lakes,  for  dyeing,  501. 

Lampblack,  277. 

Landolt,  demonstration  of  conserva- 
tion of  matter,  6. 

Langworthy,  respiration  calorimeter, 
313. 

Lanthanum,  compounds,  503. 


586 


INDEX 


Latent  heat,  former  use  of  term,  74. 

Laudanum,  343. 

Laughing  gas,  215. 

Lavoisier,  demonstration  of  the  com- 
position of  air,  19  ;  determination 
of  the  composition  of  air,  227; 
system  of  nomenclature  applied  to 
minerals,  356. 

Law,  Avogadro's,  89. 

Law,  Boyle's,  34 ;  Boyle's,  relation  to 
Avogadro's  law,  94;  Charles,  38; 
Charles,  relation  to  Avogadro's 
law,  94 ;  Dulong  and  Petit,  396. 

Law,  Faraday's,  438. 

Law  for  depression  of  freezing  points 
of  solutions,  112. 

Law,  natural,  1 ;  of  combining  vol- 
umes, 89 ;  of  combining  weights, 
13 ;  of  constant  proportions,  12 ; 
of  diffusion  of  gases,  58;  of 
gravitation,  1;  of  "mass  action," 
149 ;  of  multiple  proportions,  87 ; 
of  partial  pressures,  41,  77. 

Laws,  graphical  representation  of  gas, 
42,  43  ;  of  gases,  34,  38. 

Lead  acetate,  preparation,  uses,  519  ; 
basic,  519 ;  carbonate,  prepara- 
tion, 519  ;  basic,  manufacture,  520  ; 
comparison  with  lithopone,  470 ; 
chloride,  formation,  solubility,  518 ; 
chromate,  527 ;  chromate,  dis- 
covery of  chromium  in,  524. 

Lead  dioxide,  formation,  use  in 
storage  batteries,  516 ;  contrast 
with  barium  peroxide,  518  ;  glazes, 
danger  from,  502 ;  in  brass  and 
bronze,  431. 

Lead  monoxide,  .  manufacture,  use, 
515 ;  nitrate,  preparation,  use,  519. 

Lead,  occurrence,  metallurgy,  513 ; 
properties,  uses,  514 ;  alloys,  515 ; 
oxides,  515 ;  compounds,  518. 

Lead  oxide,  reduction  by  potassium 
cyanide,  321. 

Lead,  treatment  by  Parke's  process, 
alloys  with  zinc,  441. 

"  Lead  "  pencils,  276. 

Lead  peroxide,  515 ;  plumbate,  red 
lead,  preparation,  515 ;  structure, 
516 ;  plumbate,  structure,  535 ; 
red  oxide  of,  515;  sugar  of,  519; 
sulfate  in  storage  batteries,  516 ; 
sulfate,  preparation,  use  as  pig- 
ment, 519 ;  sulfide,  formation, 
solubility,  518;  sulfide,  theory  of 
precipitation,  169 ;  tetrachloride, 
preparation,  hydrolysis,  518 ;  tetra- 
sulfate  in  storage  batteries,  518. 

Leather,  chrome  tanning  of,  527. 

Leblanc,  discovery  of  soda  process, 
450. 


Leblanc  soda  process,  411 ;  dis- 
covery, 400 ;  recovery  of  sulfur  in, 
456. 

Le  Chatelier,  explosion  waves,  301 ; 
principle  of  van't  Hoff-,  111. 

Legislation  controling  Leblanc  soda 
process,  411,  456;  manufacture  of 
matches,  243 ;  control  of  lead  glazes, 
502 ;  to  prevent  poisoning  by  white 
lead,  521. 

Leguminous  plants,  fixation  of  nitro- 
gen by,  199. 

Length,  unit  of,  31. 

Levulose,  see  fructose. 

Lewis,  equilibrium  between  silver 
oxide,  silver  and  oxygen,  443 ; 
relation  of  law  of  Dulong  and  Petit 
to  Avogadro's  law,  397. 

Liebermann,  discovery  of  structure  of 
alizarin,  341. 

Life  of  an  element,  474. 

Lignites,  brown  and  black,  composi- 
tion, 280. 

Ligroin,  289. 

Lime  kilns,  continuous  and  inter- 
mittent, 452;  slaking  of,  453. 

Lime,  manufacture,  452 ;  theory  of 
formation  from  calcium  carbonate, 
453  ;  in  minerals,  357. 

Lime-nitrogen,  463. 

Lime-sulfur  wash,  164. 

Lime,  used  to  prepare  absolute 
alcohol,  325. 

Linde,  liquid  air  machine,  234. 

Liquid  air,  232;  preparation  of 
oxygen  from,  20. 

"  Liquid  smoke,  "  contains  acetic 
acid,  329. 

Liter,  defined,  31. 

Litharge,  reduction  with  blowpipe, 
304;  separation  of  lead  from 
silver  by  formation  of,  439. 

Lithium,  atomic  weight  of,  396 ; 
carbonate,  396 ;  chloride,  deter- 
mination of  molecular  weight  of, 
131 ;  comparison  with  magnesium, 
396;  hydride,  395;  nitride,  396; 
nitride,  formation,  201. 

Lithium,  occurrence,  properties,  395 ; 
flame  reaction,  396;  perchlorate, 
use  in  determining  the  atomic 
weight  of  chlorine,  131 ;  phos- 
phate, 396;  urate,  396. 

Lithopone,  470;  compared  with 
white  lead,  521. 

Liversidge,  gold  in  sea  water,  446. 

Lockyer,  discovery  of  helium  in  sun, 
237. 

Louisiana,  sulfur  in,  161. 

Luminosity  of  flames,  cause  of,  299. 

Luna,  alchemical  name  for  silver,  444. 


INDEX 


587 


Lunar  caustic,  444. 
Lunge,  theory  of  manufacture  of  sul- 
furic  acid  by  chamber  process,  178. 
Lutecium,  rare  earth  metal,  506. 
Luteocobalt  chloride,  559. 
Luteorhodium  chloride,  563. 

H,  definition,  262. 

iu/x,  defined,  262. 

Mabery,  announcement  of  Cowles 
furnace,  495. 

McCay,  formation  of  arsenic  penta- 
sulfide,  261. 

McCoy,  atoms  of  metallic  elements, 
94. 

Maclnnes,  table  for  degree  of  ioni- 
zation,  383,  384. 

Magnalium,  497. 

Magnesia  usta,  479. 

Magnesite,  478. 

Magnesium  ammonium  arsenate, 
259 ;  ammonium  chloride,  an- 
hydrous magnesium  chloride  from, 
480;  ammonium  phosphate,  use, 
decomposition  to  magnesium  pyro- 
phosphate, 481 ;  carbonate  in  hard 
water,  463. 

Magnesium  chloride,  by-product  in 
ammonia  soda  process,  413 ;  ef- 
fect on  the  formation  of  gypsum, 
460 ;  hydrate,  conduct  on  heating, 
480 ;  in  salt,  405. 

Magnesium  compounds,  effect  of 
ammonium  hydroxide  on  solutions 
of,  491 ;  diammonium  phosphate, 
decomposition,  252  ;  exercises,  493  ; 
hydroxide,  formation,  theory  of 
solubility  in  solutions  of  ammo- 
nium salts,  479,  491 ;  metaphos- 
phate,  formation,  252. 

Magnesium  nitride,  formation,  201 ; 
occurrence,  preparation,  proper- 
ties, 478 ;  uses,  479  ;  oxide,  prepa- 
ration, uses,  479. 

Magnesium  pyrophosphate,  forma- 
tion, 252 ;  pyrophosphate,  from 
magnesium  ammonium  phosphate, 
481;  silicide,  349;  sulfide,  hy- 
drolysis, 491 ;  sulfide,  preparation, 
hydrolysis,  480. 

Magnetic  oxide  of  iron,  formation 
from  iron  and  steam,  49 ;  from 
burning  iron,  23 ;  ore,  structure, 
556. 

Magnetite,  540. 

Malachite,  428;  formation,  431. 

Malonic  acid,  carbon  suboxide  from, 
316. 

Maltodextrin,  from  starch,  334. 

Maltose,  alcohol  from,  325 ;  forma- 
tion from  starch,  hydrolysis,  334. 


Mammoth  cave,  saltpeter  from,  418. 

Manchot,  combination  of  ferrous  sul- 
fate  with  nitric  oxide,  217 ;  ferrous 
chloride  and  nitric  oxide,  554. 

Manganates,  preparation,  536. 

Manganese  dioxide,  history  of  uses, 
535 ;  use  in  preparing  oxygen,  21. 

Manganese  heptoxide,  properties, 
538 ;  list  of  oxides,  538 ;  occur- 
rence, properties,  533  ;  alloys,  uses, 
compounds,  534 ;  valence,  structure 
of  compounds,  534 ;  tetrachloride, 
probable  formation,  101. 

Manganic  acid,  change  to  perman- 
ganic acid,  536. 

Manganous  chloride,  530. 

Manganous  hydroxide,  535. 

Manganous  manganic  oxide,  534, 
structure,  535. 

Manganous  sulfides,  535. 

Manometer,  41. 

Maple  sugar,  333. 

Marsh's  test  for  arsenic,  257. 

"  Mass  action,"  law  of,  149. 

Mass  and  weight,  relation,  32 ;  de- 
pendent on  velocity,  5. 

Matches,  242  ;  from  tetraphosphorus 
trisulfide,  243 ;  law  forbidding 
ordinary  phosphorus  in,  243. 

Matte,  copper,  429  ;   nickel,  559. 

Matter,  conservation  of,  6 ;  defini- 
tion, 5 ;  indestructibility  of,  6. 

Mauve,  discovery  of,  340. 

Mechanical  energy,  units  of,  33 ; 
equivalent  of  heat,  6. 

Medicine,  relation  of  radiochemistry 
to,  475. 

Meerschaum,  trisilicate,  356. 

Meker  burner,  temperature  of,  303. 

Melting  point,  criterion  of  pure  sub- 
stance, 12. 

Melting  points  of  elements,  372, 
table,  373 ;  in  absolute  tempera- 
ture, 135. 

Mendeleef,  identification  of  scan- 
dium as  ekaboron,  503 ;  gallium  as 
eka-aluminium,  506 ;  periodic  sys- 
tem, 136. 

Mental  work  in  respiration  calorim- 
eter, 316. 

Menzies,  critical  temperature  of  mer- 
cury, 485. 

Mercuric  chloride,  ionization  of 
anomalous,  382;  chloride,  prep- 
aration, 489;  properties,  uses, 
antidote  for,  490;  cyanide,  prep- 
aration, decomposition,  490 ;  de- 
rivatives of  ammonia,  492. 

Mercuric  fulminate,  use,  hydrolysis 
to  hydroxylamine,  491 ;  iodide,  for- 
mation, complex  compound  with 


588 


INDEX 


potassium  iodide,  use,  490  ;  iodide, 
Nessler's  reagent  from,  492 ;  ni- 
trate, 490. 

Mercuric  oxide,  decomposition,  9 ; 
formation  and  decomposition,  19 ; 
preparation,  488  ;  yellow,  489. 

Mercuric  sulfide,  red  and  black,  use, 
489 ;  solubility,  491. 

Mercurous  chloride,  formation  in 
amalgamation  process,  441 ;  chlo- 
ride, preparation,  vapor  density, 
formula,  uses,  489 ;  nitrate,  prep- 
aration, oxidation  and  reduction 
of,  490;  basic,  490;  oxide,  488; 
sulfate,  use  in  Weston  and  Clark 
cells,  437. 

Mercury,  abnormal  ionization  of 
compounds  of,  491 ;  exercises,  493  ; 
heat  of  combustion,  27 ;  occur- 
rence, 484;  metallurgy,  purifica- 
tion, properties,  critical  tempera- 
ture, 485 ;  uses,  amalgams,  486  ; 
use  in  recovering  gold,  446 ;  use 
in  Castner-Kellner  process,  402. 

Metaboric  acid,  366. 

Metachloroantimonates,  268. 

Metallic  elements  in  periodic  system, 
136. 

Metallurgy,  development  of,  390 ; 
electrolytic  methods  in,  391 ;  of 
aluminium,  history,  391;  roasting 
of  sulfides  in,  391 ;  use  of  fuels  in, 
390. 

Metals,  characteristics,  369. 

Metals,  classification,  370,  371; 
preparation  by  thermite  process, 
495  ;  spectra  of,  427  ;  systematic 
study  of,  390. 

Metantimonic  acid,  267. 

Metaphosphoric  acid,  formation 
properties,  253;  hydrolysis,  249; 
properties,  polymeric  forms,  254. 

Metasilicates,  355. 

Metasilicic  acid,  355. 

Metastannic  acid,  preparation,  prop- 
erties, 512. 

Metastannyl  chloride,  512. 

Metathesis,  defined,  81. 

Meteorites,  iron  and  nickel  in,  540. 

Meter,  definition,  31. 

Methane,  kindling  temperature,  288. 

Methane,  limits  for  explosive  mix- 
ture with  air,  288 ;  occurrence, 
preparation  from  sodium  acetate, 
properties,  286;  structure,  284; 
substitution  products  of,  287. 

Methyl  alcohol,  manufacture,  324, 
properties,  use,  325;  relation  to 
structure  of  methyl  ether,  324. 

Methyl  amine,  339. 

Methyl  ammonium  iodide,  339. 


Methyl  chloride,  from  methane,  287. 

Methylene  chloride,  from  methane, 
287. 

Methyl  ether,  determination  of  struc- 
ture, 324. 

Methyl  iodide,  relation  to  structure  of 
methyl  ether,  324. 

Methyl  red,  use  in  determining  free 
carbonic  acid,  465. 

Methyl  silicate,  colloidal  silicic  acid 
from,  353. 

Meyer,  Lothar,  preparation  of  hy- 
driodic  acid,  145. 

Meyer,  V.,  kindling  temperature  of 
methane,  288. 

Mho,  defined,  380. 

Mica,  348 ;    orthosilicate,  355. 

Michael,  explosion  waves,  301. 

Michelson,  explosion  waves,  301. 

Michigan,  bromine  from  brines  in, 
140. 

Microcosmic  salt,  use,  423  ;  formula, 
decomposition,  use,  253. 

Micron,  defined,  262. 

Migration  of  ions  in  electrolysis,  113. 

Milk  sugar,  source,  use,  334. 

Millikan,  number  of  molecules  in  Ice., 
16,  96. 

Milliliter,  used  instead  of  cc.,  31. 

Mineral,  calculation  of  formula  of, 
356. 

Mirrors,  tin  amalgam  for,  487. 

Mispickel,  256. 

Mixer  for  iron,  542. 

Mixtures  and  pure  substances,  7. 

Moissan,  artificial  diamonds,  274 ; 
preparation  of  fluorine,  153. 

Moisture,  determination  of  in  air 
by  weighing,  and  dew  point,  232 ; 
effect  on  chemical  reactions,  312 ; 
precipitation  from  air  at  an  alti- 
tude, 232  ;  presence  in  air,  231 

Mol,  definition,  183. 

Molar  solutions,  definition  of,  183. 

Molecular  weights,  determined  by 
measuring  osmotic  pressure,  360. 

Molecules  of  the  elements,  93; 
number  of  in  1  cc.,  95,  16. 

Molybdenite,  528. 

Molybdenum,  occurrence,  properties, 
528 ;  compounds,  529  ;  trioxide, 
properties,  528;  use  in  molybdic 
solution,  529. 

Molybdic  anhydride,  complex  com- 
pounds from,  529. 

"  Molybdic  solution  "  preparation, 
use  to  determine  phosphoric  acid, 
529,  530. 

Molybdic  sulfate,  reduction  of  molyb- 
dic acid  to  for  phosphorus  de- 
terminations, 529. 


INDEX 


589 


Monatomic  gases,  ratio  of  specific 
heats,  236. 

Monazite  sand,  cerium  from,  363, 
thorium  from,  364. 

Mo  no  calcium  phosphate,  249  ;  solu- 
bility, hydrolysis,  461. 

Monoclinic  system  (crystallography), 
195. 

Monopotassium  diarsenite,  259. 

Monosodium  phosphate,  249  ;  methyl 
orange  as  indicator  for,  251. 

Montana,  arsenic  from  smelting  fur- 
nace in,  256. 

Moore,  dissociation  and  ionization 
of  ammonium  hydroxide,  204. 

Mordants,  342  ;  potassium  pyrochro- 
mate,  527  ;  titanium  compounds  as, 
363. 

Morley,  determination  of  the  compo- 
sition of  water  by  volume,  68 ;  by 
weighing  oxygen  and  hydrogen, 

Morphine,  343. 

Mortar,  composition,  hardening  of, 
454. 

Mother  of  vinegar,  329. 

Muffle  furnace,  440. 

Multiple-effect  evaporators,  405. 

Multiple  proportion,  law  of,  87. 

Munroe-Neubauer  crucibles,  565. 

Muscular  energy  in  respiration  calo- 
rimeter, 315 ;  346. 

Musgrave,  development  of  Leblanc 
soda  process,  400. 

Naphtha,  289. 

Naphthalene,  from  coal  tar,  use,  295. 

"  Nascent,"  definition,  213. 

Natrium,  11. 

Natural  gas,  286. 

Natural  law,  definition,  1. 

Nature  of   chemical  energy,  27 ;    of 

scientific  knowledge,  1. 
Negatives  in  photography,  445. 
Neodymium,    discovery,    separation 

from   praseodymium,    compounds, 

504. 

Neon,  discovery,  238. 
Nernst,  equilibrium  for  formation  of 

nitric  oxide,  216  ;  quantum  theory, 

398  ;  table  for  periodic  system,  135 ; 

lamp,  363. 

Nessler's  reagent,  490,  492. 
Neubauer-Munroe  crucibles,  565. 
Neutrality,  definition  of,  385. 
Neutralization,    definition,    121 ;    in 

liquid  ammonia  solutions  by  union 

of  hydrogen  and  amide  ions,  208 ; 

theory  of  imperfect,  386,  384. 
New  Caledonia,  nickel  from,  559. 
Newton's    law  of    inertia,    probable 


basis  for  principle  of  van't  Hoff-Le 
Chatelier,  111. 

Nickel  carbonyl,  preparation,  proper- 
ties, 561 ;  chloride,  560 ;  dimethyl- 
glyoxime,  separation  from  cobalt, 
560  ;  occurrence,  properties,  559  ; 
uses,  alloys,  compounds,  560 ;  sul- 
fate,  560. 

Nicotine,  342. 

Niobium,  see  Columbium. 

Niton,  discovery,  238;  properties, 
474 ;  half-life  of,  474. 

Nitrates,  formation  in  soils,  199 : 
oxides  prepared  from,  392. 

Nitric  acid,  action  on  copper,  213; 
action  on  zinc,  213 ;  and  nitric 
oxide  from  nitrogen  peroxide,  220  ; 
as  oxidizing  agent,  212 ;  decom- 
position of,  212 ;  detection  with 
ferrous  sulfate,  217 ;  formation  of 
nitrates  from,  212. 

Nitric  acid,  hydrates  of,  211 ;  oxida- 
tion of  feathers  or  wool  by,  212 ; 
preparation  from  sodium  nitrate, 
210 ;  properties,  212 ;  structure 
according  to  electron  theory,  207. 

Nitric  oxide  and  carbon  bisulfide, 
flame  of,  217  ;  and  ferrous  chloride, 
554 ;  and  nitric  acid  from  nitro- 
gen dioxide,  220  ;  coefficient  of  ex- 
pansion, 38 ;  combination  with 
ferrous  sulfate,  217 ;  deviation 
from  Boyle's  law,  35 ;  equilibrium 
of  formation,  35;  formation  of 
nitrites  from,  217 ;  formed  by  re- 
duction of  nitric  acid  with  ferrous 
sulfate,  217 ;  formed  by  union  of 
nitrogen  and  oxygen,  215 ;  prepa- 
ration by  action  of  nitric  acid  on 
copper,  215 ;  structure,  217 ;  use 
in  manufacture  of  sulfuric  acid, 
178. 

Nitrobenzene,  use  in  preparing  ani- 
line, 340. 

Nitrogen,  coefficient  of  expansion,  38 ; 
combination  with  oxygen,  200 ; 
with  hydrogen,  201 ;  deviation 
from  Boyle's  law,  35 ;  exercises, 
225;  iodide,  preparation,  proper- 
ties, 225;  liquid,  234;  list  of 
oxides  of,  214 ;  pentoxide,  prepa- 
ration, properties,  220 ;  dioxide, 
formation  from  nitric  oxide,  219  ; 
dioxide,  from  copper  nitrate,  434 ; 
dioxide,  from  nitric  oxide,  217; 
ratio  of  specific  heats  of,  237 ;  tri- 
chloride, endothermic,  225 ;  tri- 
chloride, equivalent  to  3  Ch,  224 ; 
trichloride,  formation  by  action  of 
chlorine  on  ammonia,  209  ;  occur- 
rence and  natural  history,  198; 


590 


INDEX 


trichloride,  oxidation  of  arsenious 
oxide  by,  224 ;  trichloride,  prepara- 
tion, properties,  224 ;  peroxide, 
structure,  219 ;  peroxide,  use  in 
manufacture  of  sulfuric  acid,  178 ; 
preparation  from  air,  200  ;  prepara- 
tion from  ammonium  nitrite,  200  ; 
from  sodium  nitrite,  200 ;  proper- 
ties, 200 ;  sources  of  for  vegetable 
growth,  198 ;  tetroxide,  dissocia- 
tion, 219  ;  tetrqxide,  formation  of 
nitrous  and  nitric  acids  from,  219 ; 
tetroxide,  relation  to  nitrogen  per- 
oxide, 219;  tetroxide,  structure,  219. 

Nitroglycerin,  326  ;  explosion  of,  327. 

Nitro  Nitrogen  Trichloride,  probable 
existence,  225  ;  hydrolysis  of,  226. 

Nitrosyl  chloride,  properties  and 
hydrolysis,  214. 

Nitrosylsulfuric  acid,  formation  in 
manufacture  of  sulfuric  acid,  178. 

Nitrous  acid,  formation,  218 ;  anhy- 
dride, dissociation,  218 ;  anhydride, 
preparation,  218  ;  anhydride,  use  in 
manufacture  of  sulfuric  acid,  178 ; 
oxide,  deviation  from  Boyle's  law, 
35 ;  oxide,  preparation,  properties, 
structure,  uses,  214 ;  oxide,  use  as 
anesthetic,  215. 

Nomenclature,  Lavoisier's  system 
applied  to  minerals,  356 ;  of  acids 
and  salts,  123 ;  of  binary  com- 
pounds, 29  ;  of  oxides,  29. 

Noncoking  coals,  281. 

Nonmetallic  elements  in  periodic 
system,  136. 

Non-metals,  characteristics,  369. 

Non-sol  glass,  467. 

Nordhausen  sulfuric  acid,  556. 

Normal  solutions,  definition  of,  184. 

Norton,  composition  of  silicic  acids, 
354. 

Norway,  manufacture  of  nitrates  in, 
460. 

Noyes,  A.  A.,  cause  of  slight  solu- 
bility of  cobalt  and  nickel  sulfides, 
558. 

Noyes,  W.  A.,  determination  of  the 
composition  of  water,  71. 

Number  of  molecules  in  1  cc.,  16. 

Nutrition,  diet,  345,  347;  studied 
with  respiration  calorimeter,  345. 

Octahedron  (crystal),  193. 
Octane,  283. 
Octene,  283. 
Octine,  283. 

Octovalent,  definition,  64. 
Ohm  defined,  33. 

Oil  gas,  296 ;  percentage  composi- 
tion, 299. 


Oil  of  vitriol,  46, 

Oleic  acid,  source,  331. 

Olein,  332. 

Opal,  composition,  354. 

Open  hearth  steel,  548 ;  regenerative 
furnaces  for,  549,  550. 

Opium,  343. 

Ordinates,  axis  of,  43. 

Ores,  assay  of  for  gold  and  silver, 
440. 

Origin,  mathematical  definition,  43. 

Orpiment,  256 ;  preparation,  proper- 
ties, uses,  260. 

Orthoantimonic  acid,  267. 

Orthoclase,  trisilicate,  356. 

Orthophosphoric  acid,  classes  of  salts, 
249 ;  decomposition  of  salts  of, 
252 ;  formation,  248 ;  ionizatipn, 
250  ;  preparation,  249  ;  properties, 
249;  solubility  of  salts  of,  252; 
structure,  247. 

Orthosilicates,  355. 

Orthosilicic  acid,  355. 

Oscillators,  in  quantum  theory,  398. 

"  Osmic  acid,"  see  osmium  tetroxide, 
565. 

Osmium,  catalyzer  for  synthesis  of 
ammonia,  201. 

Osmium-iridium,  composition,  prop- 
erties, 564 ;  ruthenium  in,  563. 

Osmium,  occurrence,  oxides,  564 ; 
chlorides,  osmates,  565;  tetroxide, 
use  in  histology,  565. 

Osmosis,  358. 

Osmotic  pressure,  connection  with 
freezing  points  and  boiling  points 
of  solutions,  360;  defined,  360; 
measurement  of,  359. 

-ous,  suffix,  meaning,  30 ;  use  for 
acids,  123. 

Oxalic  acid  and  hydrazine  from  bis- 
diazoacetic  acid,  222 ;  carbon  mo- 
noxide from,  311;  decomposition, 
466 ;  ionization,  relation  to  solu- 
bility of  calcium  oxalate,  465,  466  ; 
occurrence,  manufacture  from  so- 
dium formate,  329  ;  strength  illus- 
trated, 386 ;  use,  decomposition, 
330. 

Oxidation,  definition,  63 ;  reactions 
for  potassium  permanganate,  537 ; 
writing  equations  for  reactions  of, 
171. 

Oxide  of  mercury,  decomposition,  9. 

Oxides,  nomenclature  of,  29 ;  of 
nitrogen,  summary  of  methods  of 
preparation,  214 ;  preparation  from 
metals,  nitrates,  carbonates  and 
hydroxides,  392 ;  by  precipita- 
tion, 393  ;  valence  of  elements  in, 
157. 


INDEX 


591 


Oxidizing  flame,  blowpipe  and  Bun- 
sen  burner,  304. 

Oximes,  222. 

"  Oxone,"  preparation  of  oxygen 
from,  21. 

Oxyacids  of  chlorine,  structure, 
130. 

Oxygen,  19 ;  absorbed  by  molten 
silver,  443 ;  and  acid  properties, 
23 ;  and  chlorine,  comparison  of 
heats  of  combination,  108 ;  basis 
of  unit  for  atomic  weights,  68 ; 
coefficient  of  expansion,  38 ;  com- 
bination with  hydrogen  reversible, 
372. 

Oxygen,  determination  of  in  air,  227  ; 
deviation  from  Boyle's  law,  35 ; 
dissociation  pressure  of  from 
barium  peroxide  and  manufacture, 
469  ;  for  medicinal  use  from  liquid 
air,  235  ;  liquid,  234  ;  occurrence, 
19 ;  of  air,  may  have  come  from 
carbon  dioxide,  230 ;  origin  of 
name,  23. 

Oxygen,  preparation  from  liquid  air, 
20  ;  mercuric  oxide,  19  ;  "  oxone," 
21 ;  potassium  chlorate,  20  ;  potas- 
sium chlorate  with  manganese 
dioxide,  21 ;  sodium  peroxide,  21. 

Oxygen,  properties  of,  22 ;  weight 
of  one  liter  in  different  latitudes, 
22. 

Oxyhydrogen  blowpipe,  61. 

Ozone,  action  on  silver,  443 ;  prepa- 
ration, 97  ;  properties,  97  ;  struc- 
ture, 98 ;  from  action  of  fluorine 
on  water,  154. 

Palladium,  semipermeable  mem- 
brane for  hydrogen,  358,  360; 
catalytic  effect,  use  as  catalyzer 
for  reduction  of  fats,  564 ;  di- 
chloride,  use  in  gas  analysis,  564 ; 
occurrence,  properties,  absorption 
of  hydrogen,  563 ;  oxides,  chlo- 
rides, ammines,  564. 

Palmaer,  absolute  potential  of  ele- 
ments, 436. 

Palmitic  acid,  source,  331. 

Palmitin,  331. 

Paper,  manufacture,  337;  sizing, 
338. 

Paraffin,  290. 

Parastannic  acid,  512. 

Paregoric,  343. 

Paris  green,  259. 

Parke's  process  for  silver,  440. 

Parsons'  apparatus  for  hydrogen 
sulfide,  165. 

Partial  pressures,  law  of,  41,  77. 

Paste,  glass,  467. 


Pattison's  process  for  silver,  439. 

Peat,  composition,  280. 

Pectin,  relation  to  jelly,  337. 

Pectose,  relation  to  jelly,  337. 

Pens,  iridium  for  tips  of,  565. 

Pentane,  283. 

Pentathionic  acid,  188. 

Pentine,  283. 

Pepsin,  344. 

Per-,  prefix,  meaning,  30,  123. 

Perchlorates,  128. 

Perchloric  acid,  128 ;  structure, 
130;  structure  of  hydrated,  129. 

Perchloric  anhydride,  129. 

Periodic  law,  exceptions  to,  138. 

Periodic  system,  132;  tables,  134, 
135. 

Perkin,  discovery  of  mauve,  340. 

Perkin,  fireproofing  of  cotton  goods, 
513. 

Permanent  hardness,  311. 

Permanganates,  preparation,  537. 

Permanganic  acid,  formed  by  use  of 
sodium  bismuthate,  269 ;  from 
manganic  acid,  537. 

Permanganic  anhydride,  538. 

Permonosulfuric  acid,  188. 

Peroxides,  structure,  518. 

Perrin,  estimate  of  number  of  mole- 
cules in  1  cc.,  96. 

Persulfuric  acid,  preparation,  uses, 
187. 

Pertitanic  acid,  use  in  detecting 
titanium,  362. 

Petrolatum,  290. 

Petroleum  ether,  289. 

Petroleum,  occurrence,  localities, 
varieties,  refining,  289. 

Pewter,  509. 

Pharmaceutical  extracts,  use  of  al- 
cohol in,  325. 

Phase,  effect  of  escape  of  a  gaseous, 
376;  solid,  effect  of  on  reaction, 
377. 

Phases,  definition,  77;  of  chlorine 
hydrate,  107. 

Phase  rule,  107 ;  dissociation  of 
calcium  carbonate,  453  ;  plaster  of 
paris,  458 ;  transition  or  quadruple 
point  for  sodium  sulfate,  406. 

Phenacetine,  340. 

Phenol,  source,  manufacture,  prop- 
erties, use  as  antiseptic,  326. 

Phenolphthalein,  use  as  indicator  in 
liquid  ammonia,  208 ;  use  in 
determining  free  carbonic  acid, 
464. 

Phenylhydrazine,  derivative  of  hydra- 
zine,  223. 

Phosgene,  316. 

Phosphate  rock,  241. 


592 


INDEX 


Phosphine,  compared  with  ammonia, 
arsine  and  stibine,  244;  prepara- 
tion, properties,  243. 

Phosphomolybdic  acids,  529. 

Phosphonium  group,  unstable,  244 ; 
iodide,  interference  in  preparing 
hydriodic  acid,  145 ;  iodide,  prep- 
aration, 243. 

Phosphor  bronze,  431. 

Phosphoric  acid,  determination  with 
molybdic  solution,  529,  530; 
formed  by  burning  phosphine, 
243 ;  water  soluble,  citrate-solu- 
ble, insoluble,  in  fertilizers,  461. 

Phosphorous  acid,  preparation,  prop- 
erties, 248 ;  structure,  247. 

Phosphorus,  acids  of,  list,  247 ; 
acids  of,  basicity,  247 ;  allotropic 
forms,  241 ;  burning  in  oxygen, 
23;  chlorides  of,  hydrolysis,  245; 
exercises,  255. 

Phosphorus,  heat  of  combustion,  27 ; 
in  proteins,  343  ;  occurrence,  240  ; 
oxides  of,  246;  oxychloride, 
formed  by  hydrolysis  of  the  penta- 
chloride,  245;  oxychloride,  prep- 
aration from  the  trichloride,  246. 

Phosphorus  pentachloride,  action  on 
hydroxyl  compounds,  245 ;  dis- 
sociation, 245;  hydrolysis,  116, 
245 ;  preparation,  properties,  244. 

Phosphorus  pentasulfide,  properties, 
uses,  254. 

Phosphorus  pentoxide,  efficiency  as 
drying  agent,  246 ;  from  burning 
phosphorus,  23 ;  moisture  left 
in  gas  by,  54 ;  preparation,  prop- 
erties, 246. 

Phosphorus,  positive  and  negative 
valences,  248 ;  preparation,  241 ; 
red,  241 ;  sulfides  of,  254 ;  sulfide, 
use  for  matches,  243. 

Phosphorus  tetroxide,  246 ;  tet- 
roxide, hydrolysis,  254 ;  trichlo- 
ride, hydrolysis,  1 15  ;  trichloride, 
hydrolysis,  245 ;  trichloride,  prep- 
aration, properties,  244 ;  triox- 
ide,  preparation,  properties,  246; 
valence  in  acids,  248 ;  yellow,  241. 

Phosphotungstic  acid,  531. 

Photographic  plate,  effect  of  Rontgen 
and  Becquerel  rays  on,  471. 

Photography,  444,  dry  plates,  posi- 
tives, negatives,  developing,  fixing, 
toning,  445. 

Photometry,  stellar,  use  of  selenium 
in,  190. 

Photosphere  of  sun,  spectrum,  426. 

Physical  sciences,  4. 

Physics,  definition,  5. 

Pictet,  liquefaction  of  air,  233. 


Pig  iron,  continuous  casting  ma- 
chines for,  composition,  gray, 
white,  chilled,  543;  analyses,  544. 

Pintsch  gas,  296. 

Pitchblende,  uranium  in,  531. 

Plank,  quantum  theory,  398. 

Plaster  of  Paris  in  cement,  454 ; 
manufacture,  use,  457 ;  phase 
rule,  458. 

Plating,  silver,  442. 

Platinic  chloride,  preparation,  prop- 
erties, 566. 

Platinized  asbestos,  catalysis  of 
union  of  O  and  H  by,  62 ;  prep- 
aration, 62. 

Platinous  chloride,  565. 

Platinum  ammines,  566 ;  catalyzer 
for  sulfur  dioxide,  175 ;  disulfide, 
566;  iridium  electrodes  in  alkali 
manufacture,  401 ;  metals,  gen- 
eral properties,  table,  562 ;  prints 
in  photography,  445;  properties, 
uses,  catalytic  action,  sponge, 
565. 

Plticker  tubes  for  spectra  of  gases, 
427. 

Plumbic  acid,  516. 

Plumbum,  11. 

Poisoning  by  white  lead,  521 ;  by 
lead  water  pipes,  514. 

Polarimeter,  used  to  determine  sugar, 
333 

Polarization  of  light  by  crystals,  196. 

Polarized  light,  effect  of  sugar  on, 
333. 

Poly  sulfides,  ammonium,  422. 

Polythionic  acids,  188. 

Porcelain,  501;    glazing,  502. 

Positives  in  photography,   445. 

Potassium  aluminium  sulfate,  500 ; 
argenticyanide,  320 ;  use  in  silver 
plating,  321 ;  aurate,  450 ;  aurous 
cyanide,  formation  in  cyanide 
process,  446 ;  bicarbonate,  prep- 
aration, use,  420. 

Potassium  carbonate  from  wood 
ashes,  414  ;  from  beet  sugar  manu- 
facture, from  wool,  419 ;  prop- 
erties, 419 ;  uses,  420. 

Potassium  chlorate,  composition,  21 ; 
manufacture,  uses,  416 ;  prepa- 
ration, 127 ;  preparation  of  oxygen 
from,  20. 

Potassium  chloroaurate,  450. 

Potassium  chloride,  measurement 
of  degree  of  ionization  of,  380 ; 
conductance  of  solutions  of,  380 ; 
properties,  use  in  fertilizers,  416 ; 
for  manufacture  of  saltpeter,  416, 
418 ;  chloroplatinate,  use  in  de- 
termining atomic  weight  of  chlo- 


INDEX 


593 


rine,  130;  chloroplatinate,  566; 
chloroplatinite,  use  in  photogra- 
phy, 565 ;  chloroplumbate,  519  ; 
chromate,  preparation  from  chrome 
iron  ore,  527 ;  chromium  sulfate, 
527  ;  cobaltinitrite,  formation,  prop- 
erties, use  in  analysis,  558,  419 ; 
colbalticyanide,  558 ;  cobaltocyan- 
ide,  558 ;  cupric  chloride,  use  in 
iron  analysis,  432 ;  cuprocyanide, 
435 ;  cyanate,  321 ;  cyanide,  for- 
mation, preparation,  use,  319  ;  cyan- 
ide, preparation,  use,  420. 

Potassium  diuranate,  532;  dichro- 
mate,  mordant,  342 ;  dichromate, 
reduction  by  hydrogen  sulfide, 
171.;  ferrate,  553;  ferricyanide, 
formation  from  potassium  cya- 
nide, 320 ;  ferricyanide,  use  in 
blue-print  paper,  331 ;  ferric  ferro- 
cyanide,  321 ;  ferrocyanide,  forma- 
tion from  potassium  cyanide,  320 ; 
Prussian  blue  from,  320  ;  ferro- 
cyanide, preparation,  319 ;  flu- 
oride, acid,  155 ;  fluotantalate, 
use  in  purifying  tantalum,  524. 

Potassiumhydroxide.'preparationfrom 
the  carbonate,  415 ;  by  electroly- 
sis, properties,  415 ;  slight  effect 
on  glass,  use  in  analysis,  416. 

Potassium  iodate,  reduction  in  pre- 
paring the  iodide,  417 ;  iodide, 
manufacture,  uses,  417. 

Potassium  manganate,  preparation, 
conversion  to  permanganate,  536 ; 
mercuric  iodide,  490 ;  metachlo- 
roantimonate,  268 ;  metallic,  dis- 
covery, preparation,  properties,  415. 

Potassium  nitrate,  formation  in 
soil,  199 ;  nitrate,  relation  to 
gunpowder,  417 ;  sources,  manu- 
facture, properties,  uses,  418 ; 
nitrite,  preparation,  properties, 
use,  419. 

Potassium,  occurrence,  relation  to 
minerals,  clays,  soils,  plant  growth, 
414  ;  osmate,  565  ;  oxide,  415  ; 
perchlorate,  preparation,  128 ;  per- 
chlorate,  preparation,  properties, 
416. 

Potassium  permanganate,  properties, 
uses,  537  ;  typical  oxidations  with 
537 ;  reaction  of  hydrochloric 
acid  with,  159 ;  purification  of 
hydrogen  by,  55. 

Potassium  perruthenate,  563 ;  per- 
sulfate,  preparation,  187 ;  polyi- 
odides,  formation,  use  in  iodimetry, 
417 ;  pyrochromate,  preparation, 
properties,  uses,  relation  to  pyro- 
sulfate,  527 ;  use  in  chrome  tanning, 


527 ;  pyrosulfate,  decomposition, 
use  in  analysis,  417. 

Potassium,  retention  in  soils  by  colloi- 
dal silicic  acids,  354 ;  ruthenate, 
563;  silver  cobaltinitrite,  559, 
419 ;  silver  cyanide,  use  in  silver 
plating,  442. 

Potassium,  source  for  shales,  clays 
and  soils,  494 ;  sulfate,  417 ; 
sulfate,  acid,  preparation,  pyro- 
sulfate from,  417 ;  sulfocarbonate, 
preparation,  317,  use,  318 ;  tar- 
trate,  acid,  330 ;  tetroxalate,  use 
as  standard  in  alkalimetry,  330. 

Potassium  thiocyanate,  carbon  oxy- 
sulfide  from,  318 ;  thiocyanate, 
preparation,  use  in  testing  for 
iron,  321 ;  triiodide,  145 ;  zincate, 
483. 

Potential,  absolute,  of  elements,  436 ; 
differences  of  in  relation  to  corro- 
sion of  iron,  550. 

Power,  unit  of,  33. 

Praseodymium,  discovery,  separa- 
tion from  neodymium,  compounds, 
504. 

Precipitation,  theory  of,  376. 

Premier  and  Schupp,  molecular  weight 
of  sulfur  vapor,  163. 

Preparation  of  compounds,  general 
methods,  372-379;  of  pure  sub- 
stances, 8. 

Pressure,  effect  of  on  a  gas,  34. 

Priestly,  analysis  of  air  by  nitric 
oxide,  230. 

Primary  salts,  249. 

Printing  in  photography,  445. 

Producer  gas,  297 ;  heat  relations 
in  manufacture,  298;  percentage 
composition,  299  ;  use  in  regenera- 
tive furnace  for  steel,  550. 

Propane,  structure,  284. 

Propene,  283;    structure,  284. 

Propine,  283. 

Propylene,  structure,  284. 

Proteins,  occurrence,  343 ;  diges- 
tion, 344. 

Prussian  blue,  320;  use  in  chrome 
green,  525. 

Prussic  acid,  319. 

Ptomaines,  343. 

Ptyalin,  344. 

Pure  substances  and  mixtures,  7 ; 
composition  of  expressed  in  mul- 
tiples of  atomic  weights,  17 ; 
distinguished  from  mixtures,  12; 
preparation  of,  8. 

Purpureocobalt  chloride,  559. 

Pyridine  solutions  with  semiper- 
meable  membrane,  358. 

Pyrite,  161. 


594 


INDEX 


Pyrite  burners,  179;  use  of  oxide  of 

iron  from,  540. 
Pyroantimonic  acid,  267. 
Pyroarsenic  acid,  259. 
Pyroboric  acid,  367. 
Pyrolusite,  533. 
Pyrophosphoric  acid,  hydrolysis,  248 ; 

preparation,  properties,  253. 
Pyrosulfates,  preparation,  186. 
Pyrosulfuric  acid,  186. 

Quadrivalent,  definition,  64. 
Quadruple  point  for  sodium  sulfate, 

406. 
Qualitative  analysis,  basis  for  groups 

of,  166 ;   definition,  66 ;   groups  of, 

166 ;   of  water,  66. 
Qualitative  synthesis  of  water,  66. 
Quantitative  analysis,  definition,  66. 
Quantitative  synthesis  of  water  by 

volume,  66. 
Quantum  theory,  398. 
Quartz,  348 ;   properties,  fused,  uses, 

352. 

Quinine,  343. 
Quinquivalent,  definition,  64. 

Radiations,  penetrating,  471 ;  kinds 
of,  472,  473. 

Radical,  definition,  47 ;  definition, 
relation  to  structure,  323. 

Radioactive  elements,  series  of,  475. 

Radiochemistry,  relation  to  geology 
and  medicine,  475. 

Radiothorium,  364. 

Radium,  an  element,  9,  473  ;  chemical 
action,  475 ;  discovery,  471 ;  dis- 
integration, 473  ;  emanation,  474 ; 
evolution  of  heat  by,  472 ;  half- 
life,  474 ;  nature  of  rays,  473 ; 
properties,  472. 

Rails,  steel,  manufacture,  548. 

Ramsay,  discovery  of  argon,  235 ; 
discovery  of  helium,  237  ;  discovery 
of  helium  from  radium,  473  ;  disso- 
ciation of  nitrous  anhydride,  218 ; 
possible  disintegration  of  atoms  by 
radium  emanation,  properties  of 
niton,  474,  475  ;  use  of  periodic  sys- 
tem in  discovery  of  noble  gases,  136. 

Rare  earth  elements,  position  in 
periodic  system,  134,  138. 

Rare  earths,  general,  502  ;  groups  of, 
503 ;  methods  of  separation,  503, 
504. 

Raschig,  theory  of  sulfuric  acid 
manufacture,  179. 

Rayleigh,  discovery  of  argon,  235. 

Rays,  a,  0,  473 ;  7,  5,  474. 

Rays,  chemical  action  of  radioactive, 
475. 


Reacting  substances,  effect  of  remov- 
ing one  on  equilibrium,  152. 

Reactions,  bimolecular  and  unimolec- 
ular,  150 ;  calculation  of  relative 
speed  of  at  equilibrium,  151 ; 
effect  of  insolubility  on,  376 ; 
effect  of  volatility  on,  374 ;  re- 
versible, 50 ;  speed  of  chemical, 
151. 

Realgar,  256;  preparation,  proper- 
ties, uses,  260. 

Reasoning,  inductive,  13. 

Reciprocal  ohms,  definition,  380." 

Red  lead,  oxide,  manufacture,  515; 
structure,  516. 

Reducing  agent,  sodium  amalgam, 
zinc  amalgam,  487  ;  flame,  blowpipe 
and  Bunsen  burner,  304. 

Reduction,  definition,  63 ;  writing 
equations  for  reactions  of,  171. 

Refining,  electrolytic,  of  copper,  429. 

Refrigeration  by  ammonia  in  ma- 
chines, 204. 

Regenerative  furnace  for  open  hearth 
steel,  549,  550. 

Regular  system  (crystallography) , 
193. 

Reid,  discovery  of  indium,  506. 

Resistanz  glass,  467. 

Resonators  in  quantum  theory,  398. 

Respiration  calorimeter,  313 ;  study 
of  nutrition  with,  345. 

Reversed  flames,  304. 

Reversible  reactions,  50;  hydro- 
chloric acid  and  oxygen,  109 ; 
hydrogen  and  iodine,  146 ;  ioniza- 
tion,  115;  salt  and  sulfuric  acid, 
119;  theoretically  all  reactions, 
372. 

Rhodium,  properties,  oxides,  chlo- 
rides, complex  salts,  563. 

Rhombic  dodecahedron  (crystal), 
193;  hexahedron  (crystal),  195; 
system  (crystallography),  194. 

Richter,  discovery  of  indium,  506. 

Roasting,  defined,  263;  sulfides  in 
metallurgy,  391. 

Rochelle  salt,  use  in  Fehling's  solu- 
tion, 335. 

Rock  crystal,  use  for  lenses,  352. 

Rock  salt,  mining  and  obtaining,  404. 

Rontgen,  discovery  of  Rontgen  rays, 
471. 

Rosa,  respiration  calorimeter,  313. 

Roscoe  and  Schorlemmer,  Treatise, 
standard  for  ventilation,  231. 

Rose,  distinguished  columbium  and 
tantalum,  523. 

Roseocobalt  chloride,  559. 

Roseorhodium  chloride,  563. 

Rouge,  555. 


INDEX 


595 


Rubidium  alum,  424 ;  chloroplatinate 
424,  393 ;  discovery,  occurrence 
in  carnallite,  properties,  424. 

Ruby,  494  ;   artificial,  500. 

Rum,  325. 

Ruthenium,  occurrence,  oxides,  chlo~ 
rides,  double  chlorides,  ruthenates, 
perruthenates,  563. 

Rutherford,  number  of  molecules  in 
1  cc.,  16,  96  ;  disintegration  theory, 
472. 

Saccharimeters,  333. 

Sackur,  quantum  theory,  398. 

Safety  lamp,  Davy. 

Safety  matches,  242. 

Sainte-Claire-Deville,  preparation  of 
aluminium,  495. 

Saleratus,  420. 

Sal  soda,  411. 

Salt  and  sulfuric  acid,  reversible  re- 
action, 374  ;  definition,  47,  121 ; 
mining  and  obtaining,  404. 

Saltpeter,  Chili,  210;  formation  in 
soil,  199. 

Salts,  nomenclature  of,  124 ;  hy- 
drolysis, general  statement,  394 ; 
list  of  insoluble,  393  ;  solubility  of, 
general  statement,  393 ;  valence 
of  elements  in,  158. 

Samarium,  discovery,  compounds, 
505. 

Sapphire,  494,  500. 

Scandium,  discovery,  same  as  eka- 
boron,  compounds,  503  ;  same  as 
ekaboron,  136. 

Scheele,  discovery  of  chlorine,  535. 

Science,  subdivisions  of,  3. 

Sciences,  abstract,  physical,  bio- 
logical, psychological,  4. 

Scientific  knowledge,  Nature  of,  1. 

Scott,  determination  of  the  composi- 
tion of  water  by  volume,  68. 

Sea  water,  gold  in,  446. 

Secondary  salts,  249. 

Selection  of  atomic  weights,  16,  92. 

Selenic  acid,  preparation,  properties, 
190. 

Selenite,  457. 

Selenium  dioxide,  decomposition  by 
sulfur  dioxide,  190 ;  preparation, 
190. 

Selenium,  occurrence,  allotropic 
forms,  uses,  189. 

Self-hardening  steels,  552. 

Semibituminous  coals,  composition, 
280. 

Semipermeable  membranes,  357 ; 
preparation,  358 ;  mechanism  of 
action,  358,  360 ;  for  pyridine, 
water,  hydrogen,  358. 


Septivalent,  definition,  64. 

Series  of  hydrocarbons,  283 ;  elec- 
tromotive, 435,  table,  436;  of 
radioactive  elements,  475. 

Serpentine,  disilicate,  356. 

Sexivalent,  definition,  64. 

Shales,  formation  of,  494. 

Sherardized  iron,  482. 

Sicily,  sulfur  in,  160. 

Siderite,  540. 

Siemens-Martin  process  for  steel, 
548 ;  regenerative  furnaces  for, 
549,  550. 

Silicates,  artificial,  352 ;  list  of  com- 
mon, 348,  355;  natural,  355. 

Silicic  acid  from  silicon  fluoride,  350 ; 
from  silicon  chloride,  351 ;  insol- 
uble after  drying,  354. 

Silicic  acids,  composition,  354 ;  im- 
portance of  colloidal  in  soils,  354 ; 
preparation,  colloidal,  dialysis  of, 
353 ;  structure  of,  354 ;  various, 
355. 

Silicon  carbide,  349. 

Silicon  dioxide,  forms,  348,  351 ; 
formula,  354 ;  properties,  352. 

Silicon,  exercises,  368;  fluoride, 
preparation,  hydrolysis,  350  ;  hexa- 
iodide,  preparation,  silicooxalic  acid 
from,  351 ;  occurrence,  348  ;  prepa- 
ration, properties,  349 ;  tetrachlo- 
ride,  preparation,  properties, 
hydrolysis,  351  ;  tetrafluoride, 
formed  in  etching  glass,  151 ; 
tetraiodide,  silicon  hexaiodide  from, 
351. 

Silicooxalic  acid,  preparation,  proper- 
ties, structure,  351. 

Silicozirconates,  363. 

Silver,  annual  production  and  value, 
properties,  alloys,  plating,  442 ; 
argenticyanide,  evidence  of  exist- 
ence of  complex  ions  in,  transfer- 
rence  of  ions  of  in  electrolysis,  379  ; 
bromide,  properties,  conduct 
toward  light,  444 ;  use  in  photog- 
raphy, 445. 

Silver  chloride,  properties,  conduct 
toward  light,  444 ;  use  in  pho- 
tography, 445  ;  reduction  by  mer- 
cury, 441;*  solubility  product  of, 
373. 

Silver  chloroplatinate,  566;  dep- 
osition from  potassium  argenti- 
cyanide, 321;  exercises,  450;  for 
mirrors,  487. 

"  Silver  from  Clay,"  applied  to 
aluminium,  495. 

Silver  hydroxide,  442  ;  ionization  of, 
443  ;  hydroxide,  temporary  forma- 
tion, 393 ;  iodide,  properties,  con- 


596 


INDEX 


duct  toward  light,  444 ;  iodide,  use 
in  photography,  445 ;  molten,  ab- 
sorption of  oxygen  by,  443. 

Silver  nitrate,  preparation,  proper- 
ties, uses,  444 ;  theory  of  reaction 
with  salt,  376 ;  use  in  determining 
the  atomic  weight  of  chlorine,  131. 

Silver  nitrite,  preparation,  use,  444 ; 
oxide,  dissociation  pressure,  443 ; 
formation,  442 ;  orthoarsenite,  259  ; 
peroxide,  formation,  443 ;  pyro- 
phosphate,  253. 

Silver,  occurrence,  metallurgy,  Patti- 
son's  process,  439 ;  cupellation, 
Parke's  process,  440 ;  amalgama- 
tion process,  electrolytic  process, 
cyanide  process,  441. 

Silver  recovered  from  copper,  430, 
439 ;  separation  from  lead,  439, 
440 ;  sulfide,  conversion  to  chloride 
in  ores,  441 ;  sulfide  on  coin,  test 
for  sulfur,  409 ;  sulfate,  prepara- 
tion, use,  444 ;  trinitride,  failure 
to  react  with  iodine  trinitride,  224. 

Sizes  for  paper,  338. 

Slag,  blast  furnace,  use  for  cement, 
454,  543. 

Slimes,  from  electrolytic  refining  of 
copper,  430. 

Smalt,  558. 

Smaltite,  557. 

Smith,  Alex.,  formula  of  calomel,  489. 

Smith,  E.  F.,  potassium  fluotantalate, 
524 ;  preparation  of  metallic  cal- 
cium, 452. 

Smithells,  separated  Bunsen  flame, 
301. 

Smithsonite,  481. 

Smokeless  powder,  419 ;  gun  cotton 
in,  338. 

Soaking  pits,  in  steel  manufacture,  548 

Soap,  manufacture,  use,  332 ;  water 
glass  in,  353 ;  soft,  from  lye  of 
wood  ashes,  414. 

Soapstone,  metasilicate,  355. 

Soda,  baking,  412;  washing,  411. 

Soda  lime,  use  in  preparing  methane, 
286 ;  to  absorb  carbon  dioxide  and 
water,  6. 

Sodamide,  base  in  liquid  ammonia, 
208 ;  hydronitric  acid  from  with 
nitrous  oxide,  223 ;  preparation, 
properties,  ionization,  use  in  mak- 
ing indigo,  410. 

Soda  water,  309. 

Soddy,  discovery  of  helium  from 
radium,  473. 

Sodium  aluminate,  hydrolysis,  499 ; 
decomposition  by  carbonic  acid, 
496 ;  aluminate,  manufacture  from 
clay,  496;  amalgams,  composi- 


tion, formula  for,  487  ;  ammonium 
phosphate  (microcosmic  salt) ,  253  ; 
antimonite,  265 ;  atoms,  rate  of 
vibration,  426. 

Sodium  bicarbonate,  formation  from 
carbonic  acid,  360 ;  loss  of  carbon 
dioxide  on  boiling  solution,  376 ; 
Solvay  or  ammonia  soda  process, 
412. 

Sodium  bismuthate,  oxidation  of 
manganese  by,  269  ;  bisulfate,  408 ; 
bisulfite,  408;  bromate,  144; 
bromate,  use  with  potassium  iodide 
to  illustrate  strength  of  acids,  386. 

Sodium  carbonate,  by-product  in 
making  aluminium  oxide  from 
clay,  496 ;  decomposition  by  acids, 
theory,  375 ;  formation  from  car- 
bonic acid,  310 ;  hydrolysis  of,  385  ; 
Leblanc  soda  process,  411 ;  mono- 
hydrate,  411 ;  transition  point, 
dekahydrate,  412 ;  uses,  412. 

Sodium  chloride,  crystallization,  con- 
centration of  solution  of,  405 ;  elec- 
trolysis of,  401 ;  localities  for,  398  ; 
properties,  essential  in  diet,  406; 
sodium  bicarbonate  from,  412 ; 
sources,  404 ;  theory  of  reaction 
with  silver  nitrate,  376. 

Sodium  cobaltinitrite,  558 ;  copper 
orthophosphate,  253 ;  formate, 
oxalic  acid  from,  329 ;  hydrosul- 
fide,  formation,  properties,  409. 

Sodium  hydroxide,  Castner-Kellner 
process  for,  402 ;  from  sodium 
carbonate,  401 ;  from  sodium 
chloride,  401 ;  properties,  density 
of  solutions,  table,  403. 

Sodium  hypobromite,  143  ;  hypophos- 
phite,  formed  in  preparing  phos- 
phine,  243 ;  hyposulfite,  prepara- 
tion, uses,  408 ;  hyposulfite,  prep- 
aration, 186 ;  hyposulfite,  use  for 
reduction  of  indigo,  341. 

Sodium  "  hyposulfite,"  old  name 
for  thiosulfate,  408. 

Sodium  hypophosphate,  acid,  254 ; 
iodide,  separation  of  ions  by  cen- 
trifugal force,  114;  manganate, 
preparation  conversion  to  perman- 
ganate, 536 ;  manufacture  for 
aluminium,  495 ;  manufacture  of 
sodium  peroxide  from,  404  ;  meta- 
phosphate,  formation,  use,  253 ; 
metaphosphate,  from  microcosmic 
salt,  use,  423. 

Sodium  nitrate,  acid  sodium  sulfate 
from  decomposition  of,  408 ;  for- 
mation in  soil,  199 ;  reduction  to 
nitrite,  218;  source,  properties, 
use,  410. 


INDEX 


597 


Sodium  nitrite,  preparation,  218; 
preparation,  use,  410. 

Sodium,  occurrence,  398 ;  metallurgy, 
properties,  399  ;  uses,  400 ;  oxide, 
preparation,  properties,  404 ;  per- 
borate, bleaching  by,  367  ;  perchlo- 
rate,  use  in  preparing  perchloric 
acid,  128. 

Sodium  peroxide,  hydrogen  peroxide 
from,  84;  preparation  of  oxygen 
from,  21 ;  preparation,  properties, 
hydrate,  uses,  404. 

Sodium  phosphites,  248 ;  pyro- 
borate,  367 ;  pyrosulfate,  prepara- 
tion, 186 ;  pyrosulfate,  prepara- 
tion, use  in  analysis,  408 ;  ses- 
quicarbonate,  occurrence,  398 ; 
silicates,  manufacture,  use,  413 ; 
silicate,  preparation,  properties, 
uses,  353  ;  sulfarsenate,  261. 

Sodium  sulfate,  acid,  preparation, 
uses,  pyrosulfate  from,  408 ;  manu- 
facture, hydrate,  transition  point  for 
hydrate,  406  ;  solubility  curve,  407  ; 
constituent  of  Hunyadi  water,  408. 

Sodium  sulfide,  preparation  from 
sodium  hydroxide,  from  sodium 
carbonate  and  sulfur,  as  test  for 
sulfur,  409 ;  sulfite,  acid,  use  in 
cider,  175 ;  sulfite,  acid,  prepara- 
tion, uses,  408. 

Sodium  stannate,  use  in  fireproofing 
cotton  goods,  513  ;  stannite,  prep- 
aration, reducing  agent,  510 ; 
tetrathionate,  from  sodium  thio- 
sulfate,  187 ;  tetrathionate,  for- 
mation in  iodine  titrations,  409. 

Sodium  thiosulfate,  conversion  to 
tetrathionate,  187 ;  preparation, 
use  in  photography,  408 ;  in 
lixiviation  processes,  409 ;  anhy- 
drous, 409  ;  precipitant  for  copper, 
409;  structure,  409;  use  in 
photography,  445. 

Sodium  trinitride,  formation,  410 ; 
trinitride,  preparation,  223  ;  tung- 
state,  531 ;  zincate,  483. 

Softening  water,  by  boiling,  463  ;  by 
use  of  sodium  salts,  464 ;  Clark's 

Erocess  for,  464  ;  "  permanently  " 
ard  waters  with  sodium  carbonate, 
phosphate,  fluoride  or  borate,  311. 

Soils,  formation  of,  494 ;  formation 
of  nitrates  in,  199 ;  importance  of 
colloidal  silicic  acids  in,  353  ;  po- 
tassium in,  414. 

Solder,  509. 

Soluble  glass,  353;  manufacture, 
uses,  413. 

Solubility  of  salts,  general  statement, 
393  ;  graphical  representation,  80. 


Solubility  product,  377 ;  rule  for,  378 ; 
relation  to  solubility  of  magnesium 
hydroxide,  479 ;  relations  for 
uni-bivalent  and  bi-bivalent  salts, 
378. 

Solute,  defined,  79. 

Solution,  defined,  79;  formular,  def- 
inition of,  183 ;  molar,  definition 
of,  183  ;  normal,  definition  of,  184  ; 
standard,  definition  of,  183. 

Solution  pressure,  defined,  435. 

Solutions,  chemical  activity  in,  81 ; 
differences  in  potential  between, 
437 ;  in  liquid  ammonia,  207 ; 
supersaturated,  80. 

Solvay,  discovery  of  ammonia  soda 
process,  400. 

Solvent,  defined,  79. 

Sommerfield,  quantum  theory,  398. 

Specific  heat,  fixes  atomic  weight  of 
indium,  506;  of  elements,  397; 
of  gases  at  constant  volume  and 
constant  pressure,  236. 

Spectra,  comparison  of,  427 ;  of 
gases,  Pliicker  tubes,  427 ;  of 
metals,  how  obtained,  427. 

Spectroscope,  direct  vision,  427 ;  de- 
scription of,  425. 

Spectrum  analysis,  424. 

Spectrum,  continuous,  bright  and 
dark  lines,  425 ;  theory  of  dark 
line,  diffraction,  solar,  426. 

Speed  of  chemical  reactions,  149 ; 
relation  to  chemical  affinity,  149  ; 
and  concentration,  149. 

Speed  relation  of  two  reactions 
calculated  at  equilibrium,  151. 

Sphalerite,  161,  481. 

Stafford,  molecular  weight  of  sulfur 
vapor,  163. 

Standard  solutions,  definition  of,  184. 

Stannate,  us£  in  fireproofing  cotton 
goods,  513. 

Stannic  acids,  table  of,  511. 

Stannic  acid,  preparation,  properties, 
511. 

Stannic  chloride,  preparation,  con- 
duct in  solution,  512  ;  oxide,  prep- 
aration, 510;  use  in  glass,  511; 
sulfide,  512.  _ 

Stannous  chloride,  preparation,  uses, 
reducing  agent,  510  ;  oxide,  510  ; 
sulfide,  510 ;  sulfide  dissolved  by 
ammonium  polysulfide,  422. 

Stannum,  11. 

Starch,  glucose  from,  334 ;  source, 
manufacture,  varieties,  335 ;  ap- 
pearance of  granules,  cooking  of, 
336. 

Stassfurt,  potassium  chloride  from, 
415. 


598 


INDEX 


Stearic  acid,  source,  331. 

Stearin,  331. 

Steel,  cementation,  analyses,  551 ; 
alloy,  552 ;  cast,  tempering,  545  ; 
theory  of  tempering,  546 ;  Besse- 
mer, 547  ;  acid  and  basic  Bessemer, 
open  hearth,  Siemens-Martin,  548  ; 
resistance  to  corrosion  increased  by 
copper,  431. 

Stellar  photometry,  use  of  selenium 
in,  190. 

Stereotype  metal,  264,  269,  515. 

Stewart,  standard  of  ventilation,  231. 

Stibine,  264. 

Stibium,  11. 

Stibnite,  263. 

Stockholm,  holmium  named  for,  505. 

Stokes,  ferric  sulfide,  556. 

Storage  batteries,  theory  of,  516. 

Storage  of  gases,  22. 

Strass,  467. 

Strength  of  acids,  definition,  168; 
illustration,  386. 

Strength  of  organic  bases,  339. 

Strong  acids  and  bases  defined,  386. 

Strontianite,  468. 

Strontium  carbonate,  468 ;  com- 
pounds, 468 ;  flame  color,  471 ; 
hydroxide,  468 ;  nitrate,  use  in 
fireworks,  468 ;  occurrence,  467  ; 
sulfate,  468. 

Structural  formulas,  basis  for,  323. 

Structure  of  compounds,  relation  to 
valence,  65 ;  of  the  oxyacids  of 
chlorine,  130. 

Strychnine,  343  ;  salt  of  chloroauric 
acid,  450. 

Strychnos  mix  vomica,  strychnine 
from,  343. 

Study  of  chemistry,  18. 

Subdivisions  of  science,  3. 

Substance,  definition,  7. 

Substantive  dyes,  342. 

Substitution,  287  ;  use  in  determining 
structure,  324. 

Sudbury,  nickel  from,  559. 

Sugar,  cane,  occurrence,  manufac- 
ture, properties,  effect  on  polarized 
light,  333;  hydrolysis,  334;  puri- 
fication by  bone  black,  278 ;  solu- 
tions clarified  by  basic  lead  acetate, 
519 ;  use  of  strontium  hydroxide 
in  manufacture  of,  468. 

Sugar  of  lead,  519. 

Sulfantimonates,  268. 

Sulfantimonites,  268. 

Sulfarsenates,  formation,  261. 

Sulfarsenites,  formation,  261. 

Sulfates,  list  of  insoluble,  183. 

Sulfides,  basis  of  groups  of  qualita- 
tive analysis,  166;  hydrolysis  of, 


171 ;  of  Group  II,  solubility,  491 ; 
roasting  in  metallurgy,  391 ;  theory 
of  precipitation  by  alkaline  sul- 
fides,  170. 

Sulfites,  as  germicides,  174;  prep- 
aration and  uses,  175. 

Sulfocarbonates,  formed  from  carbon 
bisulfide,  317. 

Sulfocarbonic  acid,  formation,  de- 
composition, 318. 

Sulfocyanates,  see  Thiocyanates. 

Sulfur,  allotropic  forms  of,  162 ; 
amorphous,  162 ;  boiling  point, 
163 ;  burning  in  oxygen,  23 ; 
by  Chance  process,  160 ;  com- 
pounds containing  halogens,  188. 

Sulfur  dioxide,  as  germicide  and 
disinfectant,  173 ;  bleaching  by, 
173 ;  catalysis  of  conversion  to 
sulfuric  acid  by  oxides  of  nitrogen, 
178 ;  coefficient  of  expansion,  38  ; 
deviation  from  Boyle's  law,  35 ; 
from  burning  sulfur,  23  ;  from  iron 
pyrites,  177  ;  preparation  by  burn- 
ing sulfur,  172 ;  preparation  by 
reduction  of  sulfuric  acid,  173 ; 
preparation  from  acid  sodium 
sulfite,  173  ;  properties,  173  ;  solu- 
tion in  water,  174 ;  uses,  173. 

Sulfur,  exercises,  196 ;  family,  table 
of  compounds,  192 ;  flowers  of, 
160;  gaseous,  Ss,  82  and  S,  163; 
group,  160 ;  heat  of  combustion,  27. 

Sulfur  hexafluoride,  188. 

Sulfur,  in  Louisiana,  getting  of,  161 ; 
in  petroleum  removed  with  copper 
oxide,  289 ;  in  proteins,  343 ;  in 
Sicily,  getting  of,  160. 

Sulfur,  lime-,  wash,  164 ;  mobile 
liquid  (Sx),  162 ;  monochloride, 
preparation,  uses,  188  ;  monoclinic, 
162 ;  native,  source  of,  160 ; 
occurrence,  160. 

Sulfur,  production  in  U.  S.  and  in 
world,  161 ;  properties,  163  ;  rhom- 
bic, 162  ;  roll  brimstone,  161 ;  test 
for  with  sodium  carbonate  on 
charcoal,  409. 

Sulfur  trioxide,  absorption  by  con- 
centrated sulfuric  acid,  176;  for- 
mation from  sulfur  and  oxygen 
reversible  reaction,  175 ;  poly- 
meric, 176 ;  preparation,  175 ; 
properties,  176 ;  sulfuric  acid 
from,  45. 

Sulfur,  uses,  164  ;  viscous  liquid  (Su), 
163. 

Sulfuric  acid  as  dehydrating  agent, 
182^  "chamber  acid,"  180; 
chamber  process  for,  177 ;  con- 
centration of,  180 ;  directions  for 


INDEX 


599 


dilution  of,  182 ;  dissociation  of, 
180 ;  electrolysis  of,  9,  47 ;  "  fum- 
ing," 176;  fuming  from  ferric  sul- 
fate,  556 ;  Gay-Lussac  tower,  179 ; 
hydrates  of,  181 ;  ionization  of,  181 ; 
moisture  left  in  gas  by,  54 ;  prep- 
aration by  "  chamber  process," 
177  ;  properties,  180 ;  reaction  with 
copper,  173  ;  reduction  by  hydroidic 
acid,  145 ;  reduction  by  hydro- 
bromic  acid,  142. 
Sulfuring  "  fruit,  164,  174. 

ulfurous  acid  as  reducing  agent, 
174 ;  formation,  174  ;  ionization  of, 
174;  "strength"  of,  174;  struc- 
ture, 174. 

ulfuryl  chloride,  hydrolysis,  188 ; 
preparation,  properties,  188. 

un's  corona,  helium  in,  237;  photo- 
sphere, spectrum,  426. 

uperphosphate,  calcium,  manufac- 
ture, use  in  fertilizers,  461. 

upersaturated  solutions,  80. 

plvite,  414. 

Symbols  of  elements,  11. 

ympathetic  ink,  557. 

Synthesis,  definition,  66. 

ynthesis,  quantitative,  of  water  by 

volume,  66. 

/stems  of  crystallography,   193. 

able  of  atomic  weights,  10 ;  abso- 
lute potentials  of  elements,  elec- 
tromotive series,  436 ;  acidity  or 
alkalinity  of  indicators,  388 ; 
compounds  of  elements  of  Group  V, 
271 ;  compounds  of  sulfur  family, 
192  ;  correction  of  apparent  weight 
of  water  to  volume,  73  ;  correction 
of  readings  of  barometer  for  glass 
and  brass  scale,  36 ;  correction  of 
readings  of  barometer  for  latitude 
and  altitude,  37 ;  degree  of 
ionization,  acids,  383 ;  bases,  383  ; 
salts,  384 ;  density  and  volume  of 
water,  73 ;  density  of  gases,  95 ; 
density  of  solutions  of  sodium  hy- 
droxide, 403 ;  deviation  of  gases 
from  Boyle's  law,  35 ;  deviation  of 
gases  from  law  of  Charles,  38 ;  dis- 
sociation of  water,  61 ;  elements  in 
earth's  crust,  11 ;  elements  of  sul- 
fur group,  160 ;  equilibrium  be- 
tween nitrogen  and  oxygen,  216; 
equilibrium  of  hydrogen,  iodine  and 
hydriodic  acid,  147;  groups  and  fam- 
ilies of  elements,  371 ;  heat  of  com- 
bustion of  C,  S,  P,  Fe  and  Hg,  27  ; 
halogen  acids,  139 ;  halogen  family, 
139  ;  nonmetallic  elements,  348 ; 
Periodic  System,  134,  135 ;  proper- 


ties and  compounds  of  elements  of 
Group  VIII,  562 ;  rate  of  decom- 
position of  hydriodic  acid,  148; 
rate  of  formation  and  decomposi- 
tion of  hydriodic  acid,  150  ;  stannic 
acids,  511 ;  values  of  calorie,  10°- 
30  °,  33  ;  vapor  pressure  of  ice  and 
water,  75 ;  vapor  pressure  of  sys- 
tems containing  calcium  sulfate, 
459  ;  varieties  of  ferrite,  546. 

Tafel,  preparation  of  hydroxylamine, 
221. 

Talc,  metasilicate,  355. 

Tannic  acid,  mordant,  342. 

Tantalum,  discovery,  properties,  523, 
use,  compounds,  524,  electric  light, 
524. 

Tartar  emetic,  266. 

Tartaric  acid,  structure,  source,  salts, 
330. 

Tasmania,  tin  from,  508. 

Teeth,  amalgam  for,  487. 

Telluric  acid,  190. 

Tellurium,  anomalous  position  in 
periodic  table,  138 ;  atomic  weight 
of,  190  ;  dioxide,  190  ;  occurrence, 
preparation,  compounds,  190. 

Tellurous  acid,  190. 

Temperature,  absolute,  39,  critical, 
233 ;  effect  of  on  a  gas,  38 ;  inter- 
national scale  of,  32  ;  kindling,  24  ; 
of  flames,  302  ;  of  interior  of  earth, 
relation  of  radioactivity  to,  476; 
units  of,  32. 

Temperatures,  thermometers  for  high, 
486. 

Temporary  hardness,  310. 

Terbium  compounds,  505. 

Terne  plate,  509. 

Tertiary  salts,  249. 

Tetragonal  pyramid  (crystal),  194. 

Tetragonal  system  (crystallography), 
194. 

Tetrahedron  (crystal),  193. 

Tetrahexahedron  (crystal),  193. 

Tetramethyl  ammonium  hydroxide, 
analogy  with  ammonium  hydrox- 
ide, 206. 

Tetraphosphorus  heptasulfide,  254. 

Tetraphosphorus  trisulfide,  proper- 
ties, 254 ;  use  for  matches,  243. 

Tetrathionic  acid,  188. 

Thallium,  discovery,  compounds,  507. 

Theorem  of  Le  Chatelier,  111. 

Theory,  atomic,  14 ;  electrochemical, 
influence  on  formulas  of  minerals, 
356;  the  electron,  181;  electron, 
in  relation  to  metals  and  non- 
metals,  370;  electron,  relation  to 
ionization,  206 ;  colloidal  solu- 
tions, 262 ;  decomposition  of  car- 


600 


INDEX 


bonates  by  acids,  375;  chamber 
process  for  manufacture  of  sulfuric 
acid,  178 ;  hydrolysis  of  salts,  385 ; 
neutralization,  384 ;  the  quantum, 
398  ;  storage  batteries,  576 ; 

Thermite  process,  Goldschmidt's, 
497  ;  chromium  by,  524 ;  tungsten 
by,  530. 

Thermodynamic  scale,  melting  points 
on,  373. 

Thermometers,  glass  for,  467 ;  of 
fused  quartz,  352 ;  mercury  and 
international  scales  for,  486 ;  zero 
point  correction,  special  for  high 
temperatures,  486 ;  point  on, 
fixed  by  transition  point  of  sodium 
sulfate,  406. 

Thio-,  prefix,  187. 

Thiocyanate,  potassium,  321. 

Thiosulfates,  preparation,  use,  186. 

Thiosulfuric  acid,  formation  and  de- 
composition, 187 ;  formation  from 
sulfur  monochloride,  188. 

Thomson,  heat  of  combustion  of 
hydrogen,  65. 

Thompson,  J.  J.,  the  electron  theory, 
181 ;  atoms  of  metallic  elements,  94. 

Thompson,  Sir  William,  size  of  mole- 
cules, 16. 

Thorianite,  364. 

Thorium,  in  monazite  sand,  proper- 
ties^ dioxide,  sulfate,  nitrate,  use 
in  Welsbach  mantles,  364;  series 
of  elements,  475. 

Thulium,  compounds,  506. 

Thyroid  gland,  iodine  in,  144. 

Tilkerode,  thallium  from,  507. 

Time,  units  of,  32. 

Tincture,  defined,  144. 

Tin,  occurrence,  sources,  metallurgy, 
508 ;  properties,  alloys,  uses,  tin 
plate,  509  ;  compounds,  510 ;  re- 
covery from  tin  scrap,  508. 

Titanium,  compounds  of  as  mordants, 
363 ;  occurrence,  properties,  com- 
pounds, 362 ;  oxide,  solution  in 
sodium  pyrosulfate,  408;  separa- 
tion from  silica,  detection,  362 ; 
tetrafluoride,  362 ;  test  for  hydro- 
gen peroxide,  86. 

Tolman,  separation  of  ions  by  cen- 
trifugal force,  114. 

Toluene,  283. 

Toning  in  photography,  445. 

Tool  steel,  high-speed,  530. 

Topaz,  349. 

Torpedoes,  gun  cotton  in,  338. 

Tourmaline,  349. 

Toxins,  344. 

Transition  or  quadruple  point,  for 
sodium  sulfate,  406. 


Transition  points,  for  steel  and  ferrite, 
546. 

Trautz,  theory  of  sulfuric  acid  manu- 
facture, 179. 

Triammonium  dodekamolybdate,  529. 

Tribasic  acids,  defined,  183. 

Tricalcium  phosphate,  249. 

Triclinic  system  (crystallography), 
196. 

Tridymite,  352. 

Trimethyl  amine,  204. 

Triphosphorus  hexasulnde,  254. 

Triple-effect  evaporation,  405. 

Triple  point,  definition,  78. 

Triple  point  water-ice-water-vapor 
above  0  °,  407. 

Trisilver  phosphate,  253. 

Trisodium  phosphate,  249 ;  aliza- 
rine green  as  indicator  for,  251. 

Trisilicates,  356. 

Trisilicic  acids,  355. 

Trisodium  phosphate,  hydrolysis,  al- 
kaline reaction  of,  252. 

Trithionic  acid,  188. 

Trivalent,  definition,  64. 

Trypsin,  344. 

Tuberculosis,  relation  to  ventilation, 
231. 

Tungsten,  history,  preparation,  prop- 
erties, use  in  lamps,  in  tool-steel, 
530. 

Tungsten  bronze  colors,  531. 

Turmeric  paper,  test  for  boric  acid,  368. 

Turnbull's  blue,  321. 

Tuyeres  of  blast  furnace,  541. 

Type  metal,  264,    175. 

Typhoid  fever  from  impure  water 
supply,  83. 

Ultramarine,  natural  and  artificial, 
502. 

Ultra-violet  light,  use  in  purifying 
water,  83. 

Uni-bivalent  salts,  law  of  solubility 
product  not  general  for,  378. 

Unimolecular  reactions,  150. 

Unit  for  atomic  weights,  68 ;  electri- 
cal charge,  438;  of  length,  31; 
of  power,  33  ;  of  temperature,  32  ; 
of  volume,  31 ;  of  weight,  31. 

Units,  absolute,  33;  electrical,  33; 
of  energy,  32;  of  heat,  33;  of 
mechanical  energy,  33  ;  of  time,  32. 

Uni-univalent  salts,  law  of  solubil- 
ity product  for,  378. 

Univalent,  definition,  64. 

Univariant,  definition,  78. 

Unsaturated  compounds,  definition, 
291. 

Uraninite,  531 ;   helium  in,  237. 

Uranium    chlorides,    532 ;    catalyzer 


INDEX 


601 


for  -synthesis  of  ammonia,  201; 
occurrence,  531,  properties,  com- 
pounds, 532 ;  radium  from,  473, 
half-life,  474 ;  series  of  radioac- 
tive elements,  475 ;  sulfate,  532. 

Uranyl  acetate,  532 ;  compounds, 
532 ;  nitrate,  532. 

Urea,  formed  in  body  and  from  am- 
monium cyanate,  345. 

Use  of  Indicators,  387  ;  for  weak  acids 
and  bases,  389. 

Vacuum  desiccator,  84;  high,  by 
means  of  charcoal,  278. 

Valence,  definition,  63;  illustra- 
tion, 63 ;  of  elements  in  oxides 
and  salts,  157 ;  of  elements  in 
periodic  system,  133 ;  relation  to 
equivalents  and  Faraday's  law, 
438 ;  use  in  determining  structure, 
323  ;  use  in  writing  equations,  156. 

Vanadinite,  522. 

Vanadium,  occurrence,  properties, 
uses,  compounds,  522. 

Vanadous  compounds,  522. 

van't  Hoflf,  definition  of  osmotic 
pressure,  360. 

van't  Hoff-Le  Chatelier,  principle  of, 
111 ;  applied  to  reversible  reac- 
tion, 148;  applied  in  the  prepara- 
tion of  sulphur  trioxide,  175 ; 
applied  in  synthesis  of  ammonia, 
201. 

Vapor  pressure,  definition,  74;  of 
hydrates,  82  ;  of  ice  and  water,  75. 

Vaselin,  290. 

Vauquelin,  discovery  of  chromium, 
524. 

Vegetable  foods,  347. 

Venetian  red,  555. 

Ventilation,  230  ;  standard  of ,  231 ;  lack 
of  causes  disease,  231. 

Vermilion,  489. 

Vinegar,  329. 

Vitriol,  blue,  433  ;  oil  of,  46, 434 ;  defi- 
nition, 434 ;  green,  554  ;  white,  483. 

Volatility,  effect  on  reactions,  374. 

Volcanoes,  source  of  carbon  dioxide, 
229. 

Volt,  defined,  33. 

Volume,  unit  of,  31. 

Washing  soda,  411. 

Water,  as  a  solvent,  79 ;  calculation 
of  the  composition  by  weight,  68 ; 
of  crystallization,  82 ;  degree  of 
ionization,  383;  density  of  at 
different  temperatures,  73. 

Water,  determination  of  composition 
by  copper  oxide,  69 ;  by  weighing 
oxygen  and  hydrogen,  71. 


Water,  determination  of,  in  air  by 
weighing  and  dew  point,  232. 

Water,  dissociation  of,  59 ;  effect  of 
on  chlorides,  112;  niters,  charcoal 
not  efficient  in,  278 ;  heat  of  fusion, 
74  ;  heat  of  vaporization,  74. 

Water  gas,  heat  relations  in  manu- 
facture, 298 ;  percentage  composi- 
tion, 299;  enriched,  297;  per- 
centage composition,  299,  296 ; 
carbon  monoxide  in,  297. 

Water  glass,  353  ;  of  hydration,  82  ; 
hydrolysis  of  chlorides  by,  115; 
ionization  of,  171,  383 ;  maximum 
density,  effect,  72 ;  phases  and 
triple  point,  78 ;  properties  of,  72  ; 
purification  of,  83. 

Water,  qualitative  analysis  and  syn- 
thesis of,  66 ;  quantitative  ioniza- 
tion of  in  relation  to  indicators,  387  ; 
sea,  amount  of  carbon  dioxide  in, 
230. 

Water,  use  of  aluminium  sulfate  in 
purifying,  500. 

Water  vapor,  effect  of,  on  the  volume 
of  a  gas,  76.  • 

Water,  vapor  pressure  of,  75 ;  table, 
75 ;  weight  of  1  liter  at  different 
temperatures,  73. 

Water-soluble  phosphoric  acid,  461. 

Waters,  effervescent,  309  ;  hard,  310 ; 
natural,  82 ;  radioactive,  mineral, 
476. 

Watt  defined,  34. 

Waves,  explosion,  301. 

Weak  acids  and  bases,  defined,  386. 

Weber,  preparation  of  chloroplatinic 
acid,  566. 

Weight  and  mass,  relation,  32 ;  unit 
of,  31. 

Weights  and  Measures,  International 
Bureau  of,  31. 

Welding  by  thermite  process,  498; 
use  of  borax  in,  367. 

Weldon  process,  103. 

Welsbach,  resolution  of  dydimium, 
504. 

Welsbach  mantles,  manufacture,  364 ; 
theory  of,  365. 

Wentzki,  theory  of  sulfuric  acid 
manufacture,  179. 

Werner,  formula  for  ammoniocupric 
sulfate,  434;  periodic  table,  138; 
theory  of  isomeric  hydrates  of 
chromic  chloride,  526 ;  theory  of 
valence,  559. 

Weston  cell,  electromotive  force,  437, 
438. 

Whisky,  325. 

White  lead,  manufacture,  properties, 
520  ;  poisoning  by,  521. 


602 


INDEX 


White  vitriol,  483. 

Willemite,  481. 

Wilson,  discovery  of  scandium,  503. 

Winkler,  discovery  of  germanium, 
361. 

Witherite,  468. 

Wohler,  discovery  of  aluminium,  495  ; 
synthetic  urea,  345. 

Wolfram,  11. 

Wolframite,  530. 

Wood  alcohol,  325;  charcoal,  277; 
oak,  composition,  280. 

Wood  spirit,  324. 

Wood's  metal,  269. 

Woolrich,  Sherardized  iron,  482. 

Wrought  iron,  development  of 
puddling  process,  544 ;  replace- 
ment by  other  irons,  545. 

Xenon,  discovery,  238. 
X-rays,  474. 
Xylene,  283. 

Yaryan  evaporator,  405. 
Yeast,  fermentation  of  sugar  by,  325. 
Ytterbium,  compounds,  504. 
Yttrium  group  of  rare  earths,   503 ; 
occurrence,  compounds,  503. 

Zechentmayer,  ferrous  chloride  and 

nitric  oxide,  554. 
Zero,  absolute,  40. 


Zero  group,  236. 

Zero  point,  corrections  for  ther- 
mometers, 486 ;  depression  of  in 
thermometers,  467. 

Zinc  chloride,  preparation,  properties, 
use  as  wood  preservative,  483 ; 
compounds,  effect  of  ammonium 
hydroxide  on  solutions  of,  491 ; 
hydroxide,  formation,  amphoteric, 
483. 

Zinc,  in  brass  and  bronze,  431 ;  oxide, 
green  with  cobalt  nitrate,  558. 

Zinc,  occurrence,  metallurgy,  proper- 
ties, uses,  galvanized  iron,  481 ; 
preparation,  properties,  uses,  482  ; 
Sherardized  iron,  482. 

Zinc  sulfate,  483 ;  use  in  Clark  cell, 
437. 

Zinc  sulfide,  fluorescence  of,  with 
radioactive  elements,  472  ;  forma- 
tion, conduct  toward  acids,  483 ; 
^  solubility,  491. 

Zinc,  use  in  Parke's  process,  alloys 
with  lead,  441.  . 

Zincic  acid,  salts  of,  483.  S 

Zirconates,  363. 

Zirconium  dioxide,  use  in  Nernst 
lamp,  oxyhydrogen  light  and  for 
crucibles,  363. 

Zirconium,  occurrence,  properties, 
compounds,  uses,  363. 

Zymase,  344. 


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